MFN 


CHEMISTRY 


Edmund  O'Neill 


ELEMENTARY   CHEMISTRY 


BY  THE  SAME  AUTHOR. 

CHEMICAL  LECTURE  EXPERIMENTS. 

With  224  Diagrams.      Crown  8vo,  $3  oo. 

A    TEXT-BOOK    OF    INORGANIC 
CHEMISTRY. 

With  146  Illustrations.     Crown  8vo,  $1.75. 

LONGMANS,    GREEN,  &   CO. 
NEW  YORK,  LONDON,   AND   BOMBAY. 


ELEMENTARY 

INORGANIC  CHEMISTRY 


BY 

G.    S.   NEWTH,   F.I.C.,   F.C.S. 

DEMONSTRATOR   IN    THE   ROYAL   COLLEGE   OF    SCIENCE,    LONDON 
ASSISTANT    EXAMINER    IN   CHEMISTRY,    SCIENCE   AND   ART    DEPARTMENT 

AUTHOR    OF 

"A    TEXT-BOOK    OF    INORGANIC    CHEMISTRY,"    "CHEMICAL 
LECTURE    EXPERIMENTS,"    ETC. 


LONGMANS,    GREEN,    AND    CO 

LONDON  AND   BOMBAY 
1899 

All  rights  reserved 


UNIFORM  WITH  THIS    VOLUME. 
Crown  8vo. 

ELEMENTARY    PHYSICS. 

By  W.  WATSON,  B.Sc.,  Demonstrator  in  Physics  in 
the  Royal  College  of  Science,  London  ;  Assistant 
Examiner  in  Physics,  Science  and  Art  Department. 

LONGMANS,   GREEN,  AND  CO. 
NEW    YORK,     LONDON,    AND    BOMBAY. 

IN  MEMOR1AM 


PREFACE 

THIS  book  has  been  written  to  meet  the  modern  and  practical 
methods  of  science  teaching  which  are  now  being  universally 
recognized  and  adopted  in  schools  and  colleges. 

Formerly  students  were  taught  chemistry  in  the  lecture- 
room,  the  knowledge  so  gained  being  supplemented  by  a 
minimum  amount  of  practical  work,  and  that  almost  ex- 
clusively analytical.  The  tendency  of  the  present  day  is  to 
make  the  student,  from  the  very  beginning,  an  investigator ; 
to  train  and  develop  his  faculties  for  observation;  to  make 
him  find  out  facts  and  discover  truths  for  himself;  in  other 
words,  to  make  him  think  instead  of  merely  committing  to 
memory  what  others  have  thought.  I  have  therefore  en- 
deavoured, as  far  as  it  is  possible  to  do  so  in  a  text-book,  to 
fall  into  line  with  these  views.  In  actual  practice  the  purely 
inductive  method  of  instruction  breaks  down.  There  is  so 
much  that  the  student  is  required  to  learn,  that  life  itself  is 
not  long  enough,  and  certainly  the  limited  time  at  the  dis- 
posal of  the  student  is  all  too  short,  to  admit  of  his  going 
through  the  necessarily  slow  process  of  gaining  this  knowledge 
by  his  own  investigations.  Some  facts  he  must  take  on  trust, 
and  the  question  therefore  resolves  itself  into  the  judicious 

889770 


vi  Preface. 

selection  on  the  part  of  the  teacher  of  the  facts  he  will 
endeavour  to  let  his  students  find  out  for  themselves,  and  those 
he  will  teach  them,  and  expect  them  to  commit  to  memory. 

In  a  text-book  it  is  almost  inevitable  that,  in  giving  such 
directions  as  will  lead  a  student  on  to  the  discovery  of  a  fact, 
the  fact  itself  shall  be  stated* 

Before  introducing  the  student  to  the  study  of  any  of  the 
elements,  I  have  sought  to  familiarize  him  with  a  number  of 
important  common  laboratory  processes,  in  a  chapter  on 
"  Simple  Manipulations ; "  and  this  is  followed  by  short 
chapters  on  the  "  Fitting  up  of  Apparatus,"  and  "  Simple 
Glass-blowing  Operations." 

After  hydrogen,  oxygen,  and  water  have  been  studied,  I 
have  introduced,  under  the  head  of  "Simple  Quantitative 
Manipulations,"  a  number  of  experiments  or  exercises  in- 
volving the  operations  of  weighing  and  measuring.  These 
experiments  have  been  selected  with  the  object  of  leading 
the  student  on  to  the  discovery  of  some  of  the  fundamental 
laws  of  chemistry,  making  use  of  such  knowledge  of  chemical 
facts  as  he  has  already  gained.  In  order  that  he  may  do 
these  with  an  entirely  unprejudiced  or  unbiassed  mind,  they 
have  intentionally  been  introduced  before  he  has  learnt  the 
use  of  symbols  and  formulae,  or  how  to  calculate  what  results 
he  ought  to  get.  For  some  of  these  experiments  I  am 
indebted  to  the  suggestions  of  Dr.  Tilden,  made  during  the 
course  of  a  short  series  of  lectures  to  science  teachers,  at  the 
Royal  College  of  Science,  London,  in  July,  1895. l 

1  Since  embodied  in  a  little  publication,  ' '  Hints  on  the  Teaching  of 
Elementary  Chemistry  "  (Longmans,  Green,  and  Co.). 


Preface.  vii 

In  an  elementary  practical  text-book  it  would  obviously 
be  unwise  to  take  up  the  study  of  more  than  a  very  limited 
number  of  the  elements  and  their  compounds.  Exactly  which 
elements  and  which  of  their  compounds  are  the  most  suitable 
for  the  purpose  is  probably  a  point  on  which  teachers  will 
hold  different  opinions.  I  have  selected  those  which,  in  my 
judgment,  are  the  best  adapted  for  an  elementary  course,  and 
which  I  consider  are  well  calculated  to  give  the  student  a 
broad,  general  foundation  upon  which  he  can  afterwards  build. 

G.  S.  N. 


TABLE   OF   CONTENTS 


I.    STATES  OF  MATTER— SOLID,  LIQUID,  GAS— PHYSICAL 

AND  CHEMICAL  CHANGES l 

II.  COMPOUNDS— ELEMENTS  — METALS— NON-METALS- 
MIXTURES— CHEMICAL  AFFINITY— CHEMICAL  AC- 
TION    8 

III.  SIMPLE  MANIPULATIONS— SOLUTION— DECANTATION— 
FILTRATION— PRECIPITATION— CRYSTALLIZATION- 
FUSION — DISTILLATION— COLLECTING  GASES  .  .  15 

IV.     FITTING  UP  APPARATUS 28 

V.     SIMPLE  GLASS-BLOWING 35 

VI.     HYDROGEN 4° 

VII.     OXYGEN — ACIDS — ALKALIES — BASES — SALTS     ...      54 

VIII.    WATER  (PEROXIDE  OF  HYDROGEN) 72 

IX.  MEASURING — WEIGHING— METRIC  SYSTEM— THERMO- 
METERS— BAROMETER — BALANCE 85 

X.    RELATION  OF  VOLUME  OF  GASES  TO  HEAT  AND  TO 

PRESSURE — THE  CRITH 97 

XI.    SIMPLE  QUANTITATIVE  MANIPULATIONS — FIRST  AND 

SECOND  LAWS  OF  CHEMICAL  COMBINATIONS    .     .     104 
XII.    HYDROCHLORIC  ACID— CHLORIDES — TESTS  FOR     .     .     120 

XIII.  CHLORINE  (THE  HALOGENS) 128 

XIV.  FURTHER  QUANTITATIVE  MATTERS — EQUIVALENCE  .     136 
XV.    ATOMIC  THEORY — ATOMS — MOLECULES 146 

XVI.     CHEMICAL  NOTATION — SYMBOLS — FORMULA — EQUA- 
TIONS       .    .     152 


xii  Table  of  Contents. 

CHAPTER  PAGE 

XVII.  QUANTITATIVE  SIGNIFICANCE  OF  EQUATIONS — COM- 
BINATIONS BY  VOLUME — GAY-LUSSAC'S  LAW — Avo- 
GADRO'S  HYPOTHESIS  —  DENSITIES  —  MOLECULAR 

WEIGHTS — UNIT  VOLUMES 158 

XVIII.    AIR — DIFFUSION — COMBUSTION 166 

XIX.    NITROGEN   AND   ITS  COMPOUNDS.      NITROGEN,  AM- 
MONIA    176 

XX.     NITROGEN  AND  ITS  COMPOUNDS  (continued}.    NITRIC 

ACID,  NITRATES 188 

XXI.     NITROGEN  AND  ITS  COMPOUNDS  (continued).    OXIDES 

OF  NITROGEN 197 

XXII.    OZONE 206 

XXIII.     CARBON 210 

XXIV.  CARBON  DIOXIDE — CARBONATES — QUANTITATIVE  DE- 
TERMINATION OF  CARBON  DIOXIDE 220 

XXV.    CARBON  MONOXIDE 231 

XXVI.     SULPHUR 236 

XXVII.    SULPHURETTED  HYDROGEN 243 

XXVIII.    SULPHUR  DIOXIDE — SULPHUROUS  ACID — SULPHITES  .  250 

XXIX.  SULPHUR  TRIOXIDE— SULPHURIC  ACID — SULPHATES  .  255 

XXX.  SOME  COMMON  CARBON  COMPOUNDS 264 

XXXI.    SIMPLE  QUALITATIVE  ANALYSIS 272 


LIST   OF   ELEMENTS 


Names. 

Atomic 
Symbols 

Atomic 
Weights 

Names. 

Atomic 
Symbols 

Atomic 
Weights. 

Aluminium  . 

Al 

27 

Molybdenum 

Mo 

96 

Antimony  (Stibium] 

Sb 

120 

Nickel      . 

Ni 

59 

Argon 

A 

(?) 

Niobrum  . 

Nb 

937 

Arsenic 

As 

75 

Nitrogen  . 

N 

H 

Barium 

Ba 

137 

Osmium    . 

Os 

191 

Beryllium  *  . 

Be 

9 

Oxygen     . 

O 

16 

Bismuth 

Bi 

207-5 

Palladium 

Pd 

106 

Boron 

B 

ii 

Phosphorus 

P 

31 

Bromine 

Br 

80 

Platinum  . 

Pt 

*95 

Cadmium     . 

Cd 

112 

Potassium    (Ka- 

Ccesium 

Cs 

133 

lium)     . 

K 

39 

Calcium 

Ca 

40 

Rhodium  . 

Rh 

104 

Carbon 

C 

12 

Rubidium 

Rb 

85 

Cerium 

Ce 

141 

Rutheni2tm 

Ru 

!°3'5 

Chlorine 

Cl 

35  '5 

Samarium 

Sm 

150 

Chromium    . 

Cr 

52 

Scandium 

Sc 

44 

Cobalt 

Co 

59 

Selenium  . 

Se 

79 

Copper  (Cuprum} 

Cu 

63 

Silicon 

Si 

28 

Didymium  . 

Di 

H5 

Silver  (Argentum) 

Ag 

1  08 

Erbium 

Er 

1  66 

Sodium(lVatrtum) 

Na 

23 

Fluorine 

F 

19 

Strontium 

Sr 

87-3 

Gallium 

Ga 

70 

Sulphur    . 

S 

32 

Germanium  .          . 

Ge 

72 

Tantalum 

Ta 

182 

Gold  (Auruni) 

Au 

197 

Tellurium 

Te 

I25 

Hydrogen    . 

H 

i 

Thallium  . 

Tl 

2037 

Indium 

In 

"3 

Thorium  . 

Th 

232 

Iodine 

I 

127 

Tin  (St  annum)  . 

Sn 

118 

Iridium 

Ir 

192-5 

Titanium- 

Ti 

48 

Iron  (Ferrttm) 

Fe 

56 

Tungsten  . 

W 

184 

Lanthanum  . 

La 

138-5 

Uranium 

U 

239-8 

Lead  (Plumbum)  . 

Pb 

207 

Vanadium 

V 

5  1'1 

Lithium 

Li 

7 

Ytterbium 

Yb 

173 

Magnesium  . 

Mg 

24 

Yttrium  . 

Y 

89-6 

Manganese   . 

Mn 

55 

Zinc 

Zn 

65 

Mercury   (Hydrar- 

Zirconium 

Zr 

90-4 

gyrum)      . 

TT 

200 

\ 

Those  printed  in  italics  may  be  regarded  as  rare  substances. 


ELEMENTARY 

PRACTICAL    CHEMISTRY. 

CHAPTER    I. 

STATES   OF   MATTER — PHYSICAL   AND   CHEMICAL   CHANGE. 

ALL  material  things  with  which  we  41  e  acquainted*  exist  under 
ordinary  circumstances  in  one  of  three  states^or  conditions — 
either  they  are  solids,  like  chalk,'  iron,:  salpWufr^  ic^  ;or  liquids, 
as  water,  alcohol,  mercury  ;  or  they  are  gases,  lik'e  air,  oxygen, 
and  steam. 

We  know  many  substances,  however,  which  can  easily  be 
made  to  pass  from  one  of  these  states  to  the  other  ;  thus  it  is 
familiar  to  all  that  when  solid  water  (that  is,  ice)  is  gently 
warmed  it  changes  into  liquid  wafer,  and  that  when  this  liquid 
is  more  strongly  heated  it  passes  off  into  steam,  which  is  water 
in  the  gaseous  or  vaporous  state.  We  also  know  that  when 
steam  is  cooled  it  changes  back  again  to  liquid  water,  and 
that  when  this  is  further  cooled  it  is  turned  into  ice,  or  solid 
water. 

Sulphur  is  an  example  of  another  common  substance  that 
can  readily  be  caused  to  change  from  one  state  to  another. 

Experiment  I. — Take  a  piece  of  sulphur  (brimstone)  and  chip 
off  a  fragment  about  the  size  of  a  pea.  (Note  that  sulphur  is  a  pale 
yellow  solid,  easily  broken,  being  very  brittle.)  Place  the  small 
piece  in  a  clean  dry  test-tube,  and  gently  heat  it  over  a  small 
Bunsen  flame,  in  the  manner  shown  in  the  figure.  Notice  that  the 

B 


States  of  Matter. 


solid  quickly  melts,  and,  if  carefully  heated,  is  changed  into  a  pale 
yellow  liquid.  Now  heat  more  strongly,  and  observe  that  the 
colour  rapidly  darkens,  and  the  liquid  presently  boils.  It  is  now 

being  changed  from  the  liquid  to 
the  gaseous  state.  As  the  gaseous 
sulphur  reaches  the  part  of  the  tube 
a  little  removed  from  the  flame,  it 
soon  becomes  cooled  again,  and 
consequently  changes  back  again 
to  the  liquid  state.  Notice  liquid 
sulphur  collecting  on  the  upper  part 
of  the  tube  in  the  form  of  small 
yellow  drops.  Continue  heating 
until  the  whole  of  the  original  frag- 
ment of  sulphur  has  disappeared 
from  the  bottom  of  the  test-tube. 
Allow  the  tube  to  cool,  and  observe 
that  the  yellow  liquid  which  had 
F  G.  i.  ,  condensed  upon  the  upper  part 

*    «  t /gradually  changes  back  to  the  solid 

state.  In  this  experiment,  therefore,  solid  sulphur  has  been 
changed , first- ;to the  liquid  &nrl  then  to  the  gaseous  state;  and 
gaseous  sulphur  has  been  allowed  to  pass  back  again,  first  to  the 
liquid  and  then  to  the  solid  condition. 

Experiment  2. — Gently  heat  a  small  quantity  of  alcohol  or 
methylated  spirit  in  a  test-tube  fitted  with  a  cork  and  bent  glass 
tube  (as  shown  in  Fig.  2),  which  is  joined  to  a  second  test-tube 
provided  with  a  cork  with  two  bent  tubes.  This  second  tube  is 
placed  in  a  glass  containing  cold  water,  to  keep  it  cool.  The  spirit 
soon  begins  to  boil,  and  to  pass  from  the  liquid  to  the  gaseous,  or 
vaporous  state.  As  this  gaseous  alcohol  passes  into  the  cooled 
tube,  however,  it  again  returns  to  the  liquid  condition,  and  will  be 
seen  collecting  at  the  bottom  of  the  tube. 

We  see,  from  these  examples,  that  with  some  things  it  is 
simply  a  question  of  whether  they  are  heated  or  whether  they 
are  cooled  that  decides  the  particular  state  they  shall  be  in. 
This  is  also  true  of  a  number  of  other  substances  ;  for  instance, 
the  metal  mercury  (quicksilver)  is  familiar  to  us  as  a  liquid,  but 
if  we  happened  to  be  living  in  the  Arctic  regions,  we  should 
know  it  as  a  hard  solid,  resembling  lead.  Again,  spectrum 
analysis  has  taught  us  that  iron,  lead,  copper,  tin,  and  many 


.    Physical  and  Chemical  Change.  3 

other  metals  which  are  familiar  to  us  as  solids,  are  present  in 
the  sun,  and  that,  owing  to  the  intense  heat,  they  exist  there 
in  the  gaseous  state,  forming  a  part  of  the  sun's  atmosphere. 
It  is  not  necessary  to  take  these  metals  to  the  sun,  however, 
to  make  them  change  their  state  from  the  solid  to  the  gaseous. 
When  we  apply  heat  to  them,  some  of  them,  such  as  lead  and 
tin,  melt  readily  enough ;  others,  like  copper  and  iron,  require 
a  much  higher  temperature  to  make  them  pass  into  the  liquid 
state ;  while  all  of  them  can  be  boiled  and  made  to  pass  into 
the  gaseous  condition  by  means  of  the  electric  furnace. 

All  solids  and  liquids  are  visible  and  tangible  ;  gases,  on 


FIG.  2. 

the  other  hand,  cannot  be  felt,  and  in  most  cases  they  are 
invisible.  It  is  quite  impossible,  by  merely  looking  at  it,  to 
tell  whether  a  glass  bottle  is  filled  with  air,  or  hydrogen,  or 
oxygen,  or  whether  it  is  entirely  empty — that  is,  vacuous ;  hence, 
when  a  liquid  substance  passes  into  the  gaseous  state,  it  usually 
disappears  altogether  from  our  sight.  For  example,  if  a 
small  quantity  cf  water  be  left  exposed  in  a  shallow  dish  or 
saucer,  we  know  that  it  gradually  diminishes  in  quantity,  and 
finally  disappears  entirely.  In  common  language,  we  say  that 
it  has  dried  up ;  in  more  scientific  phraseology,  we  speak  of  the 
process  as  evaporation — that  is  to  say,  the  water  has  passed 


4  States  of  Matter. 

from  the  liquid  to  the  vaporous  or  gaseous  state,  in  which 
condition  it  is  invisible,  and  mingles  with  the  other  equally 
invisible  gases  of  the  air.  The  visible  cloud  which  appears 
when  steam  is  allowed  to  escape  from  a  boiler  or  locomotive 
engine,  and  which  is  popularly  called  "steam,"  consists  in 
reality  of  minute  drops  of  liquid  water,  and  is  not  water  in  the 
vaporous  or  gaseous  state.  The  true  steam  is  invisible,  but 
on  suddenly  coming  into  contact  with  cool  air  the  gaseous 
water  changes  to  liquid  water  and  becomes  visible.  That 
gaseous  water  is  invisible  may  be  proved  by  the  following 
experiment : — 

Experiment  3. — Place  a  small  quantity  of  water  in  a  large 
glass  flask,  which  is  provided  with 
a  cork  carrying  a  short  glass 
tube  bent  at  right  angles.  Boil 
the  water  rapidly,  until  a  jet  of 
steam  escapes  from  the  tube  and 
produces  the  familiar  cloud  of 
"steam."  It  is  evident  that  the 
flask  must  now  be  full  of  gaseous 
water,  but  it  is  perfectly  clear  and 
invisible.  Now  hold  immediately 
beneath  the  jet  of  steam  another 
Bunsen  flame  (as  at  2,  Fig.  3),  and 
observe  that  the  cloud  instantly 
disappears.  Although  the  same 
amount  of  steam  is  escaping  from 
the  tube,  it  is  now  invisible.  This 
is  because  the  flame  warms  the  air 
in  the  immediate  neighbourhood 
of  the  jet,  so  that  the  gaseous 
water  does  not  become  suddenly 

cooled  on  issuing  into  the  air,  and  therefore  does  not  condense  to 

the  liquid  state. 

Some  solid  forms  of  matter,  when  heated,  change  into  the 
gaseous  state  without  first  becoming  liquid  ;  that  is  to  say, 
they  pass  at  once  into  gases  without  melting.  For  example — 

Experiment  4. — Heat  in  a  dry  test-tube  a  fragment  of  ammo- 
nium chloride  (sal  ammoniac}  about  the  size  of  a  grain  of  wheat. 


FIG.  3. 


Physical  and  Chemical  Change.  5 

Notice  that  the  white  solid  does  not  melt,  but  at  once  passes  into 
vapour.  The  vapour,  on  reaching  the  cooler  parts  of  the  test-tube, 
changes  back  again  to  the  solid  state,  and  collects  as  a  white 
coating  upon  the  glass.  This  process  is  termed  sublimation;  the 
ammonium  chloride  is  said  to  sublime  when  heated.  Any  impurities 
present  in  the  original  solid,  which  do  not  change  into  vapour  at 
the  temperature  employed,  will  obviously  be  left  behind,  hence  this 
process  may  be  used  to  purify  substances  like  ammonium  chloride. 

Many  kinds  of  matter,  when  experimented  upon  in  the 
same  way  as  the  sulphur  (Exp.  i)  and  the  ammonium  chloride 
(Exp.  4),  undergo  a  different  kind  of  change,  a  change  which 
is  more  subtle  and  less  simple.  This  will  be  seen  by  the 
following  examples  : — 

Experiment  5.— Heat  in  a  dry  test-tube  a  small  quantity  of 
potassium  chlorate,  and  carefully  notice  what  takes  place.  The 
solid  quickly  melts  and  changes  into  the  liquid  state,  and  presently 
appears  to  be  boiling.  So  far  it  seems  to  behave  like  the  sulphur 
(Exp.  i).  But  is  the  liquid  potassium  chlorate  being  changed  into 
the  gaseous  state  ?  In  the  first  place,  it  will  be  noticed  that 
practically  nothing  condenses  upon  the  upper  and  cooler  part  of 
the  tube  ;  this  seems  to  imply  that  the  substance  is  not  passing 
into  the  gaseous  condition.  Apply  a  lighted  taper  to  the  mouth 
of  the  tube  ;  no  gas  is  escaping  which  will  take  fire,  but  notice  that 
the  flame  of  the  taper  becomes  brighter.  Dip  into  the  test-tube  a 
match  or  splinter  of  wood  which  has  been  lighted  and  blown  out, 
and  has  still  a  glowing  spark  upon  it ;  the  wood  will  be  rekindled 
and  burst  into  flame.  This  proves  that  something  gaseous  is 
coming  from  the  boiling  liquid.  Now  fit  a  cork  and  bent  tube 
into  the  test-tube,  and  connect  it  to  a  second  tube  arranged  as  in 
Fig.  2.  It  will  be  found  that  nothing  visible,  either  liquid  or  solid, 
collects  in  the  cold  tube.  As  the  heating  of  the  potassium  chlorate 
is  continued,  it  will  be  seen  that  the  liquid  becomes  less  fluid,  and 
finally  goes  solid  again.  By  this  experiment,  therefore,  we  find  that 
when  solid  potassium  chlorate  is  heated  it  first  becomes  liquid,  and 
then  is  changed  into  two  different  things,  namely,  a  colourless  gas 
which  does  not  change  either  to  a  liquid  or  a  solid  when  cooled 
by  cold  water,  and  a  solid  which  is  evidently  different  from  the 
original  one,  because  it  is  very  much  more  difficult  to  melt. 

Experiment  6. — Place  a  few  grains  of  mercuric  oxide  in  a  test- 
tube,  and  apply  heat.  In  this  case  the  red  solid  becomes  darker 
in  colour,  but  does  not  melt.  Gradually,  however,  there  will  collect 


6  States  of  Matter. 

upon  the  cooler  part  of  the  tube  a  sublimate  which  has  the  appear- 
ance of  a  bright  white  silvery  metal,  and  upon  dipping  into  the 
test-tube  a  glowing  splint  of  wood,  we  shall  obtain  the  same  result 
as  in  Exp.  5.  By  continuing  the  heat,  the  whole  of  the  original 
substance  will  disappear.  We  learn  from  this  experiment,  there- 
fore, that  when  mercuric  oxide  is  heated,  it  also  changes  into  two 
different  things  ;  one  of  them  being  a  colourless  gas  which  passes 
away  (presumably  the  same  gas  as  was  given  out  by  the  potassium 
chlorate  in  Exp.  5),  while  the  other  is  a  white  metal.  Examine 
this  metallic  sublimate  carefully,  and  notice  that  it  consists  of 
minute  globules  of  liquid  metal.  This  must  be  mercury,  because 
no  other  metal  is  liquid  at  the  ordinary  temperature. 

Certain  differences  between  the  kind  of  change  undergone 
by  the  mercuric  oxide  and  the  potassium  chlorate,  and  that 
experienced  by  sulphur  and  by  ice,  when  these  things  are 
heated,  will  be  evident,  (i)  Neither  the  sulphur  nor  the  ice 
change  into  two  different  kinds  of  matter  at  once.  (2)  When 
sulphur  or  ice  are  changed  by  heat  into  the  liquid  or  gaseous 
state,  the  change  is  only  a  temporary  one ;  when  cooled  again, 
they  change  back  to  their  original  condition.  On  the  other 
hand,  the  mercuric  oxide  and  the  potassium  chlorate  each 
change  into  two  different  forms  of  matter,  and  the  change 
in  each  case  is  permanent.  Changes  like  those  experienced 
by  mercuric  oxide  and  potassium  chlorate  are  called  chemical 
changes,  whilst  those  which  the  ice  and  the  sulphur  under- 
went are  distinguished  as  physical  changes.  Many  chemical 
changes  are  constantly  going  on  around  us  in  the  familiar 
processes  of  everyday  life  :  thus,  when  a  candle  burns,  the 
solid  wax  is  transformed  into  certain  invisible  gases  which 
mix  with  the  air,  and  never  return  again  to  the  original 
condition  of  the  wax;  the  candle  undergoes  a  chemical 
change.  In  our  ordinary  fires,  coal  is  converted  into  smoke 
and  certain  invisible  gases,  which  escape  into  and  pollute  the 
atmosphere,  and  a  small  residue  of  a  greyish  ash  is  left ;  the 
change  is  permanent,  and  is  a  chemical  change. 

When  an  egg  is  cooked,  the  clear  and  almost  colourless 
liquid  albumen  (white  of  egg)  is  converted  into  a  white  solid, 
which  does  not  again  return  to  the  liquid  state  on  cooling; 
the  albumen  undergoes  a  chemical  change. 


Physical  and  Chemical  Change.  j 

Human  beings  eat  bread,  meat,  vegetables,  etc. ;  these 
foods  undergo  chemical  changes,  whereby  they  are  converted 
into  flesh  and  bones,  into  invisible  gases  which  leave  the 
body  in  the  breath,  and  into  waste  products  which  leave 
the  body  in  the  perspiration  and  evacuations. 

Chemical  and  physical  changes,  however,  are  very  closely 
associated,  and  although  we  may  have  a  physical  change 
without  any  chemical  change,  all  chemical  changes  are  ac- 
companied by  a  physical  change ;  and  in  many  cases  the 
accompanying  physical  change  is  the  only  outward  indication 
which  we  have  that  a  chemical  change  has  taken  place  at  all. 
For  example — 

Experiment  7. — Take  a  glass  tube  about  two  feet  long  and 
half  an  inch  wide,  and  close  one  end  like  the  bottom  of  a  test-tube. 
Fit  a  cork  into  the  other  end,  and  slip  an  indiarubber  ring  upon 
the  tube  a  short  distance  from  the  corked  end.  About  half  fill 
the  tube  with  strong  sulphuric  acid,  and  then  gently  fill  the  tube 
up  to  the  ring  with  cold  water.  The  water,  being  much  lighter 
than  the  acid,  will  float  upon  it  without  mixing,  if  poured  in  gently. 
Now  tightly  cork  the  tube,  and  tip  it  up  and  down  two  or  three 
times  in  order  to  mix  the  contents  together.  A  chemical  change 
takes  place,  but  nothing  is  to  be  seen  :  it  will  soon  be  found,  how- 
ever, that  the  tube  is  getting  so  hot  that  it  can  scarcely  be  held  in 
the  hand.  Now  hold  the  tube  in  an  upright  position,  and  notice 
that  the  liquid  has  shrunk,  for  it  no  longer  reaches  to  the  mark 
upon  the  tube.  Here  we  see  the  chemical  change  is  accompanied 
by  a  change  of  temperature,  and  a  change  of  volume. 


CHAPTER   II. 

COMPOUNDS,    ELEMEN  L'S — METALS,     METALLOIDS — MIXTURES — 
CHEMICAL  AFFINITY,  CHEMICAL   ACTION. 

Elements  and  Compounds. — When  different  kinds  of  matter 
are  experimented  upon,  it  is  found  that  from  some  of  them 
it  is  possible  to  obtain  two  or  more  substances  totally  unlike 
the  original  matter,  while  from  others  it  is  impossible  to  obtain 
anything  essentially  different  by  any  process  at  present  known. 
For  example,  in  Exp.  5  we  found  that  from  potassium  chlorate 
we  were  able  to  obtain  a  gas  which  caused  a  glowing  splint 
of  wood  to  re-light  (a  gas  called  oxygen),  and  also  a  white 
solid  residue  which  required  a  much  stronger  heat  to  melt  it 
than  the  original  substance  did.  Again,  in  Exp.  6,  from  the 
red  mercuric  oxide  we  obtained  the  same  gas  oxygen,  and  a 
quantity  of  metallic  mercury,  two  things  entirely  different 
from  the  original.  Substances  like  potassium  chlorate  and 
mercuric  oxide  are  called  compounds.  Substances  from  which 
we  are  unable  to  obtain  anything  essentially  different,  are 
distinguished  as  elements. 

There  are  only  about  70  substances  which  are  believed 
to  be  elementary  bodies;  and  nearly  the  half  of  these  may 
be  considered  as  rareties.  The  following  list  of  thirty  includes 
all  the  most  important  of  the  elements  (for  the  complete  list 
see  inside  the  cover)  : — 

Aluminium,  Al.  Calcium,  Ca.  Gold,  Au. 

Antimony,  Sb.  Carbon,  C.  Hydrogen,  H. 

Arsenic,  As.  Chlorine,  Cl.  Iodine,  I. 

Bismuth,  Bi  Copper,  Cu.  Iron,  Fe. 

Bromine,  Br.  Fluorine,  F.  Lead,  Pb. 


Mechanical  Mixtures. 


Magnesium,  Mg. 
Manganese,  Mn. 
Mercury,  Hg. 
Nickel,  Ni. 
Nitrogen,  N. 

Oxygen,  O. 
Phosphorus,  P. 
Platinum,  Pt. 
Potassium,  K. 
Silicon,  Si. 

Silver,  Ag. 
Sodium,  Na. 
Sulphur,  S. 
Tin,  Sn. 
Zinc,  Zn. 

Of  these  elements  two  are  liquid  at  the  ordinary  tempera- 
ture, namely  bromine  and  mercury ;  five  are  gases,  chlorine, 
fluorine,  hydrogen,  nitrogen,  and  oxygen,  while  the  rest  are  solid. 
Most  of  the  solid  elements  are  metals,  and  the  above  list  will 
be  seen  to  contain  the  names  of  such  familiar  metals  as 
copper,  gold,  iron,  lead,  etc.  Those  of  the  solid  elements  in 
this  list  that  are  not  metals  are  arsenic,  carbon,  iodine,  phos- 
phorus, silicon,  and  sulphur.  Some  of  these  possess,  in  a 
greater  or  less  degree,  some  of  those  properties  which  are 
usually  associated  with  the  metals.  Thus  arsenic  and  silicon 
are  opaque  substances  having  the  power  of  reflecting  light 
from  their  surfaces,  a  property  usually  known  as  metallic  lustre. 
Carbon,  again,  has  the  power  of  conducting  both  heat  and 
electricity,  properties  which  are  almost  exclusively  charac- 
teristic of  metals.  The  name  metalloids  is  sometimes  given  to 
those  elements  which,  while  not  being  true  metals,  yet  closely 
resemble  them  in  some  of  their  properties. 

Mechanical  Mixtures. — When  two  elements  are  brought 
together,  either  nothing  happens,  or  else  a  chemical  change 
takes  place.  In  the  former  case  the  result  is  a  simple  mixture 
of  the  two  elements ;  in  the  latter  it  is  the  formation  of  a 
compound.  Similarly,  when  two  compounds  are  brought 
together,  if  a  chemical  change  takes  place  it  results  in  the 
formation  of  new  compounds,  whereas  if  no  chemical  change 
follows  we  only  obtain  a  mechanical  mixture  of  the  two 
compounds. 

Experiment  8.— Grind  to  powder  in  a  mortar  a  small  quantity 
of  potassium  chlorate,  and  mix  it  with  about  the  same  quantity  of 
powdered  white  sugar.  No  chemical  change  takes  place,  thereibre 
the  result  is  a  simple  mixture. 

Experiment  9. — Place  in  a  dry  test-tube  a  fragment  of  phos- 
phorus about  the  size  of  a  pea,  and  drop  upon  it  a  few  grains  of 
powdered  iodine.  A  chemical  change  at  once  takes  place,  great 
heat  is  evolved,  and  a  compound  of  phosphorus  with  iodine  results. 


io  Mechanical  Mixtures. 

[Phosphorus  is  a  substance  which  must  be  handled  with  great  care, 
as  it  very  easily  takes  fire.  It  is  always  preserved  beneath  water, 
and  when  a  small  fragment  is  required  for  experiment,  the  larger 
piece  should  be  taken  out  of  the  bottle  by  means  of  tongs,  placed 
on  a  plate  with  some  water,  and  there  cut  with  a  penknife.  The 
piece  to  be  used  must  be  quickly  wiped  with  blotting-paper,  and 
put  into  the  test-tube  with  the  tongs.  It  should  never  be  taken 
up  by  the  fingers,  as  the  warmth  of  the  hand  might  cause  it  to 
take  fire.] 

Experiment  io.— Place  in  a  dry  test-tube  a  few  drops  of 
mercury,  together  with  a  few  particles  of  powdered  iodine,  and 
gently  heat  the  tube.  A  chemical  change  follows,  and  a  siiblimate 
is  obtained  partly  red  and  partly  yellow.  This  sublimate  consists 
of  a  compound  of  mercury  and  iodine. 

Each  of  the  ingredients  in  a  simple  mixture  retains  its  own 
individual  and  characteristic  properties.  Thus,  if  the  mixture 
of  potassium  chlorate  and  sugar  (Exp.  8)  be  tasted,  both  the 
sweet  taste  of  the  sugar  and  the  peculiar  saline  taste  of  the 
potassium  chlorate  will  easily  be  perceived. 

Owing  to  the  fact  that  their  properties  are  retained,  the 
ingredients  of  a  mechanical  mixture  can  be  separated  from 
each  other  by  processes  which  are  purely  mechanical  or  physical, 
as  distinguished  from  chemical. 

Experiment  n.— Powder  some  potassium  nitrate  (nitre),  and 
mix  into  it  a  little  powdered  sulphur  (flowers  of  sulphur)  and  a 
small  quantity  of  fine  iron  filings.  The  result  is  a  mechanical 
mixture  of  these  things.  If  this  mixture  be  now  examined  through 
a  pocket  lens,  the  separate  particles  of  white  nitre,  yellow  sulphur, 
and  grey  iron  will  be  distinctly  visible,  lying  side  by  side  unchanged. 
Place  the  mixture  upon  a  sheet  of  white  paper  and  bring  a  magnet 
near  to  it ;  the  iron  in  the  mixture  will  be  attracted  by  the  magnet, 
and  may  in  this  way  be  drawn  away  and  entirely  removed.  The 
residue,  containing  now  only  the  nitre  and  sulphur,  should  be 
removed  to  a  test-tube,  a  little  water  added,  and  the  contents  of  the 
tube  warmed.  The  nitre  being  easily  dissolved  by  water  is  thus 
separated  from  the  sulphur,  which  does  not  dissolve.  If  the 
mixture  be  now  poured  upon  a  blotting-paper  filter  (see  p.  17) 
the  watery  solution  containing  the  nitre  passes  through,  while  the 
sulphur  remains  on  the  paper.  If  the  liquid  be  then  boiled  in  a 
little  dish,  the  water  will  evaporate  and  the  nitre  will  be  left  as  a 


Chemical  Affinity.  1 1 

white  solid  residue  in  the  dish.     In  this  way  the  three  ingredients 
have  been  separated  by  physical  processes. 

Chemical  Affinity. — In  a  compound  the  various  con- 
stituents stand  in  a  totally  different  relation  to  each  other  than 
is  the  case  with  mixtures ;  a  relation  which  is  much  closer, 
and  more  difficult  to  understand. 

The  elements  present  in  a  compound  are  said  to  be 
chemically  combined  with  each  other :  their  union  in  all  cases 
is  controlled  by  the  operation  of  a  particular  force,  which  is 
spoken  of  as  chemical  affinity. 

Consider  some  of  the  cases  of  chemical  combination 
already  referred  to.  In  Exp.  10  the  two  elements,  mercury 
and  iodine,  enter  into  chemical  union,  and  the  compound 
formed  is  called  mercuric  iodide.  If  this  scarlet  powder 
be  examined  by  the  most  powerful  microscope  it  is  quite 
impossible  to  see  either  the  mercury  or  the  iodine  it  contains. 
All  the  properties  belonging  to  mercury,  as  well  as  those 
belonging  to  iodine,  are  gone,  and  the  compound  is  endowed 
with  an  entirely  new  set  of  properties  which  are  peculiar  to 
itself. 

Again,  in  Exp.  6  we  learnt  that  mercuric  oxide  was  a 
compound  of  mercury  and  oxygen.  These  two  elements,  the 
liquid  metal  mercury  and  the  colourless  gas  oxygen,  when 
united  together  by  the  influence  of  chemical  affinity,  entirely 
lose  their  own  individuality,  and  take  on  altogether  new  habits 
— the  compound  is  a  brick-red  crystalline  solid,  it  possesses 
none  of  the  properties  of  either  mercury  or  oxygen,  and  these 
constituents  cannot  be  separated  again  by  any  purely  mechanical 
operations. 

The  familiar  substance  water,  is  a  compound  obtained  by 
the  chemical  union  of  two  colourless  invisible  gases ;  one  of 
them  (hydrogen)  a  gas  which  easily  burns,  and  the  other 
(oxygen)  a  gas  which  causes  ordinary  burning  things  to  burn 
more  quickly  and  brilliantly.  The  compound,  therefore,  which 
these  gases  give  rise  to  when  they  chemically  combine  has 
none  of  the  properties  of  the  ingredients,  its  properties  are 
diametrically  opposed  to  theirs— it  is  liquid,  it  does  not  burn, 


12  Chemical  Action. 

and  it  is  the  great  antidote  for  fire.  By  no  mechanical  method 
can  we  disengage  or  separate  these  constituents. 

Why  it  is  that  the  compound  produced,  when  two  such 
substances  as  hydrogen  and  oxygen  enter  into  chemical 
union,  should  have  the  particular  properties  with  which  water 
is  endowed,  no  one  knows.  Or  why  the  compound  of 
mercury  and  iodine  should  happen  to  be  red,  and  not  blue 
or  green,  we  cannot  tell.  And,  in  the  same  way,  it  is  quite 
impossible,  from  a  knowledge  of  the  properties  of  the  in- 
gredients, to  foretell  what  will  be  the  nature  of  the  compound 
they  will  give  rise  to.  For  instance,  given  a  certain  gas 
(chlorine),  with  a  greenish-yellow  colour,  a  powerful  irritating 
smell,  and  extremely  poisonous;  also  a  soft  white  metal 
(sodium),  which  if  placed  upon  the  tongue  would  take  fire ; 
now  what  sort  of  properties  are  likely  to  be  possessed  by  a 
compound  formed  by  the  chemical  union  of  these  substances  ? 
No  one  would  predict  that  the  product  would  be  the  innocent 
and  necessary  article  of  food,  common  salt ;  but  such  is  actually 
the  case. 

Chemical  Action  is  the  term  applied  to  the  actual  pro- 
cesses which  take  place  by  the  operation  of  the  force  called 
chemical  affinity.  Thus,  when  phosphorus  unites  with  iodine 
(Exp.  9),  or  when  iodine  combines  with  mercury  (Exp.  10), 
the  process  of  combination  is  termed  chemical  action,  we  say 
that  a  chemical  action  takes  place  between  the  phosphorus, 
or  the  mercury,  and  the  iodine. 

Chemical  action  does  not  take  place  promiscuously  between 
all  the  elements.  Some  are  made  to  combine  only  with  great 
difficulty;  others  will  not  combine  together  at  all  As  a 
general  rule,  those  elements  which  are  most  unlike  each  other 
combine  together  most  readily,  they  are  said  to  have  the 
greatest  affinity  for  each  other.  Neither  does  chemical  action 
take  place  in  all  cases  under  the  same  conditions  ;  thus  some- 
times it  takes  place  at  once  on  simply  bringing  the  substances 
together. 

Experiment  12. — Place  a  small  quantity  of  sodium  peroxide  in 
a  dry  test-tube,  and  pour  upon  the  powder  a  little  cold  water. 


Chemical  Action.  13 

Chemical  action  at  once  takes  place,  the  mixture  effervesces  briskly. 
Such  effervescence  signifies  that  a  gas  is  being  set  free,  as  one  of 
the  products  of  the  chemical  action.  If  a  glowing  match  or  a 
splint  of  wood  be  dipped  into  the  test-tube,  the  same  result  will 
follow  as  in  Exps.  5  and  6,  showing  that  the  gas  given  off  is 
oxygen.  See  also  Exp.  9. 

In  other  cases  it  is  necessary  to  employ  some  external 
energy  in  order  to  induce  chemical  action  to  begin.  In  a  large 
number  of  instances  the  application  of  heat  will  bring  about 
chemical  action.  Thus,  in  Exp.  10,  chemical  action  between 
the  mercury  and  iodine  was  induced  by  heating  the  mixture. 
This  may  also  be  exemplified  by  the  following  experiments. 

Experiment  13. — Make  a  small  heap  of  the  mechanical  mixture 
of  potassium  chlorate  and  sugar  (Exp.  8),  and  apply  a  lighted  match 
to  it.  Chemical  action  is  at  once  set  up,  and  rapidly  propagated 
throughout  the  heap. 

Experiment  14. — Place  a  few  fragments  of  copper  foil  or  wire 
in  a  test-tube,  and  pour  upon  them  a  small  quantity  of  strong 
sulphuric  acid.  While  cold  no  action  takes  place,  but  if  the  mixture 
be  warmed  it  will  first  become  muddy,  and  presently  give  off  a  gas 
(sulphur  dioxide),  which  has  a  choking  smell  like  that  produced  by 
burning  sulphur. 

Sometimes  chemical  action  is  brought  about  by  the 
influence  of  light. 

Experiment  15. — Brush  over  one  side  of  a  half-sheet  of  note- 
paper  with  a  dilute  solution  of  silver  nitrate,  using  a  clean  soft 
brush,  and  put  the  paper  away  in  a  drawer  until  dry.  Cut  out  of  a 
piece  of  brown  paper  some  figure,  either  a  capital  letter,  or  other 
design.  Lay  the  prepared  paper  upon  a  piece  of  wood  and  fasten 
the  brown  paper  figure  down  upon  it  with  one  or  two  drawing-pins  ; 
expose  it  to  bright  sunshine  for  a  short  time.  Where  the  prepared 
paper  has  been  exposed,  the  sunlight  will  cause  a  chemical  change 
to  take  place,  which  will  result  in  the  discolouration  of  the  paper, 
so  that  on  removing  the  brown  paper,  an  image  of  the  design  will 
be  seen  on  the  under  paper.  This  experiment  is  a  primitive 
photographic  process.  All  photography  depends  upon  the  influence 
of  light  in  promoting  chemical  action. 

In  a  number   of  cases,  chemical  action  is  only  able  to 


14  Chemical  Action. 

proceed  in  the  presence  of  a  third  substance,  which  itself 
remains  unchanged  at  the  conclusion  of  the  action,  and  which 
sometimes  needs  only  to  be  present  in  the  most  infinitesimal 
quantity.  These  cases  are  all  classed  together  under  the  head 
of  catalytic  actions,  the  third  substance  being  spoken  of  as  the 
catalytic  agent. 

All  known  instances  of  chemical  action  take  place  accord- 
ing to  one  of  three  general  modes.  These  will  be  better 
understood  after  other  matters  have  been  considered  (p.  155). 


CHAPTER    III. 

SIMPLE    MANIPULATIONS. 

Solution. — This  term  is  applied  both  to  the  act  of  dissolving, 
and  to  the  product  obtained  by  dissolving.     For  example — 

Experiment  16. — Throw  into  some  water  in  a  test-tube  a  little 
powdered  potassium  nitrate  (nitre} ;  and,  on  shaking,  it  will  soon 
be  entirely  dissolved.  We  say  that  the  potassium  nitrate  has 
undergone  solution,  and  we  term  the  resulting  liquid  a  solution  of 
nitre. 

The  liquid  in  which  a  substance  is  dissolved  is  called  the 
solvent ;   thus,  in  the  above  example,  the  solvent  was  water. 
If  such  a  solution  be  heated,  or 
even   allowed  to   stand   exposed 
in  an  open  dish,  the  solvent  will 
gradually    evaporate    (see    p.    3-), 
and     leave    the    dissolved    sub- 
stance behind. 

Experiment  17. — Pour  the  solu- 
tion of  nitre  from  Exp.  16  into  a 
porcelain  evaporating  basin,  and 
heat  gently  by  means  of  a  Bunsen 
with  a  rose  burner,  as  shown  in 
Fig.  4.  Continue  the  process  until 
the  water  is  all  driven  off,  and 
a  white  saline  residue  is  left  in  the 
dish.  This  is  called  evaporating 
to  dryness. 

Unfortunately   the  term  solu- 
tion   is    employed    without    dis- 
tinction, to  denote  two  essentially  different  processes  of  dissolv- 
ing.    This  will  be  best  understood  by  the  following  examples, 


1 6  Simple  Manipulations. 

Experiment  18. — Place  a  small  quantity  of  powdered  sodium 
carbonate  ('washing'  soda)  in  two  separate  glasses.  To  one,  add 
some  water,  which  will  dissolve  the  sodium  carbonate,  giving  a 
clear  solution.  We  have  therefore  made  a  solution  of  sodium  car- 
bonate in  water.  Into  the  second  glass  pour  some  dilute  hydro- 
chloric acid,  and  again  the  sodium  carbonate  dissolves,  and  a  clear 
solution  is  obtained.  In  this  second  case,  however,  one  striking 
difference  will  be  noticed,  which  is,  that  the  act  of  solution  of  the 
sodium  carbonate  in  hydrochloric  acid  is  attended  by  a  brisk 
effervescence.  This  means  that  some  gas  is  being  disengaged,  and 
that  a  chemical  change  is  taking  place.  Now  place  each  solution 
in  a  porcelain  dish  and  gently  evaporate  to  dryness.  In  both  dishes 
a  white  residue  will  be  left,  but  let  us  try  and  find  out  whether  they 
are  the  same  things  or  not.  One  property  of  sodium  carbonate 
has  been  exhibited  already  during  this  experiment,  namely,  that  an 
effervescence  takes  place  when  dilute  hydrochloric  acid  is  dropped 
upon  it  ;  let  us  therefore  use  this  property  as  a  test,  and  add  a  few 
drops  of  hydrochloric  acid  to  the  residue  in  each  dish.  In  the  case 
of  the  residue  from  the  watery  solution,  there  will  be  an  effer- 
vescence, so  that  we  may  infer  that  the  sodium  carbonate  did  not 
undergo  any  permanent  change  by  being  dissolved  in  water,  and 
is  left  unaltered  when  the  water  is  evaporated  away.  In  the 
other  case  no  effervescence  will  take  place,  showing  that  this  residue 
is  something  quite  different  from  the  original  substance. 

Experiment  19. — Place  a  few  scraps  of  copper  in  a  test-tube, 
and  add  water  ;  the  metal  does  not  dissolve.  Pour  away  the  water, 
and  add  a  little  strong  nitric  acid  ;  effervescence  takes  place,  a 
brown-coloured  offensive  smelling  gas  makes  its  appearance,  and 
the  copper  disappears.  We  say,  therefore,  that  we  have  made  a 
solution  of  copper  in  nitric  acid  ;  or  that  copper  is  soluble  in  nitric 
acid.  It  will  be  noticed  that  the  solution  has  a  blue  colour,  and 
if  it  be  evaporated  gently  to  dryness  a  blue  residue  will  be  left, 
which  evidently  is  not  copper.  It  is  a  compound  of  copper,  namely, 
copper  nitrate. 

It  will  be  evident  from  these  experiments  that  two  kinds 
of  solution  are  recognized — one  being  more  of  a  physical 
character,  in  which  the  substance  is  dissolved  without  apparent 
change,  and  from  which  it  is  again  deposited  in  its  original 
state  by  evaporating  the  solvent ;  while,  in  the  other  case,  the 
act  involves  a  chemical  change,  a  chemical  action  taking  place 
between  the  solvent  and  the  dissolved  substance,  giving  rise 


Decantation  ;  Filtration. 


to  new  compounds  entirely  different  from  either  of  these 
things.  This  latter  process  is  sometimes  distinguished  as 
chemical  solution. 

•\*\      The  process  of  solution  affords  an  important  method  for 
separating  substances  that  are  mixed  together.     Thus — 

Experiment  20. — Powder  a  piece  of  marble,  and  mix  it  with 
powdered  sodium  carbonate,  and  throw  a  little  of  the  mixture  into 
some  water  in  a  test-tube,  and  shake  it  up.  The  sodium  carbonate 
dissolves,  but  the  insoluble  marble  settles  to  the  bottom.  In  order 
to  separate  the  solution  from  the  sediment,  one  of  two  methods 
may  be  used,  namely,  decantalion  we  filtration. 

Experiment  21. — Deeantation.    Allow  the  marble  to  settle,  and 
carefully  pour  off  as  much  of  the  clear  liquid  as  possible  without 
disturbing  the  sediment.    Nearly  fill  the  test-tube  again  with  water 
and,  after  thoroughly  shaking,  allow  the 
marble  to  settle  once  more,  and  again 
decant  the  clear  liquid.     By  repeating 
this  process  the  marble  will  be  washed 
free  from  the  sodium  carbonate. 

Experiment  22. — Filtration.  The 
filtering  medium,  almost  exclusively 
used,  is  a  bibulous  paper  (like  blotting- 
paper),  known  as  filter-paper.  This  is 
usually  obtained  already  cut  in  circular 
pieces  of  various  sizes.  One  of  these 
pieces  is  made  into  a  cone  by  being 
folded  first  into  half,  and  then  at  right 
angles  into  half  again,  and  is  supported 
in  a  glass  funnel  of  such  a  size  that  the 
glass  will  project  slightly  above  the 
paper  (Fig.  5).  The  cone  is  then  wetted 
with  a  little  clean  water,  and  the  filter 
is  ready  for  use.  The  solution  to  be 
filtered  is  poured  into  the  cone  (without 
overflowing  the  paper),  and  the  in- 
soluble matter  is  arrested  by  the  paper,  while  the  solution  passes 
through  quite  clear.  When  the  whole  has  run  through,  the 
insoluble  portion  is  washed  free  from  any  adhering  solution  by  once 
or  twice  filling  the  filter  with  water,  and  each  time  allowing  the 
whole  of  the  wash  water  to  drain  through. 

By   conducting   such    an   operation    carefully,    the   exact 


FIG.  5. 


18 


Simple  Manipulations. 


FIG.  6. 


quantities  of  both  the  marble  and  the  sodium  carbonate 
present  in  the  mixture  can  be  ascertained.  For  this  purpose 
the  insoluble  residue  upon  the  filter  must  be  dried  and 
weighed  ;  and  the  filtrate,  together  with  all  the  wash  waters, 

must  be  evaporated  to  dryness, 
and  the  residue  (consisting  of 
the  sodium  carbonate)  also 
weighed.  It  will  be  obvious 
that  there  must  be  no  loss  of 
any  of  the  solution  during  the 
whole  operation.  When  pour- 
ing a  liquid  from  a  narrow 
vessel  like  a  test-tube,  there 
is  no  risk  of  it  spilling  by 
running  down  the  outside  of 
the  tube;  but  in  the  case  of 
wide  vessels  it  is  very  likely  to  do  so,  when  loss  of  the  solution 
would  arise.  This  risk  is  obviated  by  the  device  of  pouring 
the  liquid  down  against  a  glass  rod  held  lightly  against  the  edge 

of  the  vessel.  Figs.  6  and  7 
show  the  two  methods.  Again, 
in  order  to  prevent  splashing 
from  the  drops  of  liquid  falling 
from  the  point  of  the  funnel, 
the  latter  should  be  made  to 
touch  the  side  of  the  vessel 
placed  below,  so  that  the  drops 
may  run  down  the  surface  of 
the  glass.  To  facilitate  this,  the 
stems  of  funnels  are  usually  cut 
diagonally,  as  seen  in  Fig.  5. 

In  the  example  above  given, 

the  separation  was  made  by  the  physical  process  of  solution, 
as  the  sodium  carbonate  did  not  undergo  any  chemical  change 
during  the  process ;  but  quite  as  frequently  the  method  of 
chemical  solution  is  employed.  For  example — 


Experiment  23. — Mix    together  powdered  marble  and  white 


Precipitation.  19 

sand.  Both  substances  are  insoluble  in  water.  Add  to  some  of 
the  mixture  dilute  hydrochloric  acid  :  effervescence  takes  place, 
showing  that  a  gas  is  being  disengaged.  When  no  further  effer- 
vescence is  noticed  upon  adding  more  of  the  dilute  acid,  the 
mixture  may  be  filtered.  The  residue  upon  the  filter  is  the  sand. 
Sand  is  not  dissolved  by  hydrochloric  acid,  but  marble  is.  If  the 
filtrate  be  evaporated  to  dryness,  a  residue  wilt  be  obtained ;  but 
this  residue  is  not  marble,  because  if  a  drop  of  hydrochloric  acid 
be  added  to  it  no  effervescence  takes  place.  The  marble  and  the 
hydrochloric  acid  have  undergone  a  chemical  reaction,  resulting 
in  the  formation  of  new  substances,  namely,  a  gas  (carbon  dioxide), 
and  the  residue  (calcium  chloride). 

Precipitation. — Sometimes  when  one  clear  solution  is 
added  to  another,  the  resulting  mixture  is  no  longer  clear ; 
something  is  produced  which  makes  the  liquid  thick  or 
muddy. 

Experiment  24.— Dissolve  a  small  pinch  of  common  salt 
(sodium  chloride)  in  water  in  a  test-tube,  and  add  a  few  drops  of 
a  solution  of  silver  nitrate.  Instantly  the  mixture  becomes  milky, 
owing  to  the  separation  of  a  white  solid,  which  finally  settles  to 
the  bottom  of  the  tube.  The  silver  of  the  silver  nitrate  has  a 
strong  chemical  affinity  for  the  chlorine  of  the  sodium  chloride, 
and  the  compound  produced  when  these  unite  (namely,  silver 
chloride)  happens  to  be  insoluble  in  water,  and  therefore  the 
moment  it  is  formed  it  separates  out  as  a  solid.  This  process  is 
called  precipitation^  and  the  solid  which  is  produced  is  termed  the 
Precipitate. 

A  precipitate  may  be  separated  from  the  liquid  in  which 
it  is  suspended,  either  by  decantation  or  nitration  (see  p.  17). 

By  means  of  the  combined  processes  of  precipitation  and 
nitration,  it  is  possible  to  separate  by  chemical  means  sub- 
stances which  are  mixed  together  in  solution.  For  example — 

Experiment  25. — Take  a  crystal  of  copper  nitrate  and  one  of 
silver  nitrate,  and  dissolve  them  together  in  distilled  water  in  a 
test-tube.  The  crystal  of  the  copper  salt  will  impart  to  the  solution 
its  own  blue  colour,  but  the  liquid  will  be  clear.  Add  to  this  a 
few  drops  of  a  solution  of  common  salt.  Just  as  in  Exp.  24,  there 
is  again  a  white  precipitate  of  silver  chloride.  Add  the  salt  solution 


2O  Simple  Manipulations. 

drop  by  drop  until  no  more  of  the  white  precipitate  is  produced. 
The  whole  of  the  silver  originally  present  in  the  silver  nitrate  is 
now  united  to  chlorine,  and  is  thrown  out  of  the  solution  as  in- 
soluble silver  chloride.  Now  filter  the  mixture  (see  p.  17),  and  a 
clear  blue  liquid,  containing  all  the  copper  nitrate,  will  pass  through 
the  filter  while  the  silver  chloride  collects  on  the  filter.  This 
method  of  separating  is  the  basis  of  most  analytical  processes. 

Crystallization. — When  we  dissolve  any  substance  in 
cold  water,  and  continue  adding  more  of  the  substance,  a 
point  is  ultimately  reached  when  the  water  will  dissolve  no 
more.  The  water  is  then  said  to  be  saturated  with  that 
particular  substance.  If  such  a  cold  saturated  solution  be 
warmed,  it  is  then  able  to  dissolve  some  more  of  the  substance, 
until  again  a  point  is  reached  when  it  can  dissolve  no  more. 
When  such  a  hot  saturated  solution  is  again  cooled,  that 
quantity  of  the  dissolved  substance  which  the  hot  solution 
contained,  over  and  above  the  amount  which  the  cold  water 
could  dissolve,  is  thrown  out  of  solution,  and  in  many  cases 
it  is  deposited  in  the  form  of  crystals. 

Experiment  26. — Add  powdered  alum  in  small  quantities  at  a 
time  to  some  cold  water  in  a  test-tube,  constantly  shaking,  and 
allowing  each  little  portion  to  dissolve  before  adding  more.  Notice 
that  the  water  gradually  dissolves  each  additional  portion  of  alum 
more  slowly,  until  at  last  the  solution  is  saturated.  Now  heat  the 
solution  to  boiling,  and  notice  that  it  will  now  easily  dissolve  a 
further  considerable  quantity  of  the  alum.  Cool  the  solution  by 
dipping  the  test-tube  into  cold  water,  and  almost  immediately  a 
quantity  of  alum  will  be  thrown  out  of  solution  and  deposited  in 
the  form  of  minute  crystals.  Once  more  warm  the  solution,  and 
observe  that  these  crystals  again  dissolve.  Now  stand  the  test- 
tube  down,  and  allow  it  gradually  to  cool  by  itself,  and  then  notice 
that  again  crystals  have  been  deposited,  but  that  they  are  much 
larger,  so  that  their  particular  shape  can  easily  be  seen.  The  same 
quantity  is  deposited  in  both  cases,  but  when  quickly  cooled  the 
crystals  are  smaller. 

Different  substances  are  soluble  in  water  to  a  very  different 
extent;  thus  we  find  that  the  same  quantity  of  cold  water 
as  is  just  capable  of  dissolving  eight  parts  of  borax,  will  dis- 
solve four  times  as  much  nitre,  and  twenty  times  as  much 


Fusion.  2 1 

zinc  sulphate  (white  vitriol).  Owing  to  this  unequal  solu- 
bility of  various  substances,  the  process  of  crystallization  is 
often  made  use  of  in  order  to  separate  substances  from  each 
other ;  more  especially  with  a  view  to  removing  from  one  sub- 
stance any  admixed  impurities  which  are  soluble,  and  which, 
therefore,  cannot  be  got  rid  of  by  filtering. 

Experiment  27. — Mix  together  about  equal  quantities  of 
powdered  potassium  chlorate  and  potassium  dichromate  ;  place 
the  mixture  in  a  beaker,  and  pour  a  little  boiling  water  upon  it. 
Place  the  beaker  over  a  rose  burner  and  boil  the  solution,  and  add 
just  enough  boiling  water  to  entirely  dissolve  the  mixture.  The 
solution  now  has  the  orange-red  colour  of  the  potassium  dichro- 
mate. Remove  the  lamp,  and  after  a  few  minutes  stand  the  beaker 
in  a  basin  of  cold  water,  when  a  crop  of  crystals  is  soon  deposited. 
As  soon  as  the  solution  is  cold,  decant  the  clear  liquid  and  drain 
the  crystals.  Now  pour  a  little  cold  water  upon  the  crystals  in  the 
beaker,  and  again  drain  them.  Observe  that  the  solution  first 
decanted  and  the  wash  water  are  coloured,  but  that  the  crystals 
themselves  are  nearly  white.  Rinse  them  once  or  twice  more  with 
small  quantities  of  cold  water,  and  see  that  each  rinse-water  be- 
comes less  strongly  coloured  and  the  crystals  become  whiter.  This 
shows  that  potassium  dichromate  is  more  soluble  than  potassium 
chlorate  :  and  it  will  be  evident  that  by  repeating  the  process  (that 
is,  by  once  more  dissolving  the  white  salt  and  crystallizing  it  again) 
we  can  get  rid  of  every  trace  of  the  yellow  salt. 

Fusion  is  the  term  applied  to  the  process  of  converting 
a  solid  substance  into  the  liquid  state  by  the  application  of 
heat.  Thus  when  ice  is  warmed  it  enters  into  a  state  of  fusion, 
or,  in  other  words,  it  melts ;  and  when  lead  is  heated  it  also 
fuses  or  melts.  In  common  language  the  term  melt  is  often 
incorrectly  employed  to  denote  the  process  of  solution,  thus 
sometimes  it  is  said  that  sugar  melts  in  warm  water.  This 
confusion  is  to  be  carefully  avoided. 

When  a  substance  in  a  state  of  fusion  is  allowed  to  cool 
and  solidify,  it  very  commonly  assumes  crystalline  shapes  ; 
thus,  when  water  passes  into  the  solid  state  we  obtain  a 
crystalline  mass  of  ice.  The  shape  of  the  crystals  of  ice  is 
readily  seen  by  allowing  a  single  snow-flake  to  fall  upon  the 
coat  sleeve,  and  looking  at  it  through  a  pocket  lens.  The 


22  Simple  Manipulations. 

magnificent  fern-like  crystals  of  ice  upon  a  window-pane  in 
winter  are  familiar  to  all,  and  very  often  single  star-shaped 
crystals  of  great  beauty  are  to  be  seen. 

Chemical  action  often  takes  place  between  substances  in  a 
state  of  fusion }  which  is  incapable  of  taking  place  when  they 
are  only  in  solution ;  for  example — 

Experiment  28. — Dissolve  a  small  piece  of  potassium  hydroxide 
(caustic  potash}  in  water,  and  add  to  the  colourless  solution  a  few 
grains  of  powdered  manganese  dioxide.  The  black  powder  simply 
falls  to  the  bottom  of  the  solution,  and  no  action  takes  place. 
Now  place  another  piece  of  potassium  hydroxide  in  a  dry  test-tube 
and  heat  it :  the  solid  melts  to  a  colourless  liquid.  While  it  is  in 
this  ftised  condition  add  a  few  particles  of  the  manganese  dioxide, 
and  notice  a  very  different  result.  The  liquid  turns  to  a  blue-green 
colour,  and  the  black  oxide  of  manganese  has  disappeared.  If 
this  be  allowed  to  cool,  it  again  solidifies  to  a  green  solid  mass  ; 
and  when  quite  cold,  if  water  be  added,  it  dissolves,  giving  a  green 
solution.  A  chemical  action  has  taken  place  between  the  man- 
ganese dioxide  and  the  ftised  potash  forming  a  compound  called 
potassium  manganate,  but  the  solution  of  potash  was  incapable  of 
bringing  about  this  reaction. 

Distillation. — The  principle  of  this  process  has  already 
been  explained  (p.  2) ;  but  the  process  is  more  conveniently 
carried  on  by  means  of  the  apparatus  in  Fig.  8.  The 
flask,  a,  known  as  a  "  Wurtz "  flask,  contains  the  liquid  to 
be  distilled.  Its  brancli  tube  passes  through  a  cork  in  the  end 
of  the  Liebigs  condenser.  This  consists  simply  of  a  straight 
tube,  T,  which  is  jacketed  by  a  wider  tube,  W,  through  which 
a  stream  of  cold  water  circulates,  the  water  entering  at  the 
lower  branch  tube,  B,  and  passing  out  at  B'.  In  this  way  the 
inner  tube  is  kept  cold.  The  neck  of  the  Wurtz  flask  is 
fitted  with  a  cork,  through  which  a  thermometer  is  fixed,  in 
order  to  tell  the  temperature  at  which  the  liquid  distils. 

Experiment  29. — About  half  fill  the  Wurtz  flask  with  water, 
which  has  been  made  dirty  by  the  addition  of  a  small  fragment  01 
clay,  or  a  little  ink,  and  replace  the  cork  carrying  the  thermometer, 
the  bulb  of  which  must  not  reach  down  into  the  water.  Boil 
the  water  by  means  of  a  Bunsen  lamp,  and  place  a  clean  dry  flask 


Distillation.  23 

to  receive  the  distillate.  During  the  whole  experiment  a  stream  of 
cold  water  must  be  circulating  through  the  condenser.  As  soon  as 
the  water  in  the  flask  begins  to  boil,  carefully  read  the  thermometer. 


FIG.  8. 

Notice  that  the  mercury  gradually  rises  until  a  point  is  reached 
when  it  remains  stationary.  Note  this  temperature.  Observe  also 
that  the  distillate  is  perfectly  colourless,  all  the  impurity  which 
rendered  the  water  in  the  flask  dirty  remains  behind,  and  only 
clean  water  passes  over.  The  process  of  distillation  therefore 
purifies  the  water  ;  and  distilled  water  is  purer  than  ordinary 
water. 

Distillation,  also,  often  enables  chemists  to  identify  a  liquid. 
For  instance,  a  colourless  liquid  is  given  you,  which  looks  like 
water;  if  it  be  submitted  to  distillation,  and  the  temperature 
noted  at  which  the  mercury  in  the  thermometer  is  stationary 
while  the  liquid  is  briskly  boiling,  you  could  tell  at  once 
whether  it  was  water  or  not 

The  process  of  distillation  is  useful  also  as  a  method  of 
separating  liquids  which  are  mixed,  but  which  boil  at  different 
temperatures. 

Experiment  30. — Place  in  the  Wurtz  flask  a  mixture  of  about 
two  parts  of  water  and  one  of  alcohol,  and  have  ready  three  small 
clean  flasks  to  receive  the  distillate  in.  When  this  mixture  is 
distilled,  observe  that  the  mercury  in  the  thermometer  quickly  rises 


Simple  Manipulations. 


to  a  certain  point  and  then  remains  steady  for  a  time.  Presently 
it  begins  again  to  rise  ;  then  change  the  receiver  and  collect  what 
passes  over  separately.  At  length  the  mercury  is  again  steady, 
indicating  the  temperature  at  which  the  water  in  Exp.  29  distilled ; 
at  this  point  exchange  the  second  receiver  for  the  third,  and  continue 
the  process  a  little  longer. 

In  receiver  No.  i  the  liquid  consists  mainly  of  alcohol,  but 
mixed  with  a  little  water.  Pour  out  a  little  of  the  liquid  and  set  fire 
to  it,  thus  proving  that  it  is  fairly  strong  alcohol.  Receiver  No.  2 
contains  a  mixture  of  alcohol  and  a  large  proportion  of  water. 
This  liquid  will  not  burn  at  all.  The  third  receiver  contains  water 
free  from  alcohol. 

By  this  process,  therefore,  we  have  obtained  a  portion  of  one  of 
the  liquids,  namely,  the  water,  free  from  the  other  liquid,  but  have 
not  completely  separated  each  from  the  other 

Collecting  Gases. — Gases  are  usually  collected  by  causing 
them  to  bubble  through  water  into  a  bottle  or  cylinder  filled 

with  water,  and  standing 
mouth  downward  in  a 
trough  or  basin  of  water. 
This  method  is  usually 
described  as  collecting  over 
wafer,  and  the  basin  used 
is  called  a  pneumatic 
trough. 

Experiment  31. — Fill  a 
glass  cylinder  to  the  brim 
with  water,  and  slide  on  to 
the  ground  lip  a  ground 
glass  plate  so  that  no  air 
bubble  is  included.  Grasp 
the  cylinder  with  the  left  hand,  as  shown  in  Fig.  9,  and  hold 
the  glass  cover  in  its  position  with  the  forefinger  of  the  right 
hand.  Then  invert  the  cylinder,  when  it  will  be  in  the  position 
shown  in  Fig.  10,  and  lower  its  mouth  beneath  the  surface  of 
the  water  in  the  basin  or  trough,  as  in  Fig.  ir,  and  now  with- 
draw the  plate.  The  cylinder  then  remains  rilled  with  water.  In 
the  figure  a  glass  basin  is  shown,  this  is  in  order  to  render  the 
position  of  the  cylinder  visible  ;  in  the  laboratory  it  is  more  usual 
to  use  pneumatic  troughs  made  of  metal,  having  a  movable  metal 


FIG.  9. 


Collecting  Gases.  25 

shelf  with  an  oblong  hole  in  it  through  which  the  bent  end  of  a 

tube  can  be  introduced.     Take  a  piece  of  glass  tube,  bent  as  shown 

in    Fig.    12,    and    gently 

blow  through  it  so  that 

the    air  from  the   lungs 

shall  bubble  up  into  the 

cylinder.       As     the    gas 

ascends    in     bubbles,    it 

gradually    displaces    the 

water,   until   at  last  the 

cylinder  is    full    of   gas. 

To  remove  the  cylinder 

in  order  to  examine  the 

gas,  slip  the  ground  glass 

plate  beneath  its  mouth, 

and.  keep  it  in  its  place 

with  the  finger  while  the 

cylinder  is  being  turned 

over.      Now  remove  the 

plate,  and  lower  a  lighted- 

candle  on   a  wire,   or    a 


FIG-  10. 


burning  taper  bent  as  shown  in  Fig.  13,  into  the  gas  in  the  cylinder. 
Notice  that  the  flame  is  extinguished,  showing  that  the  gas  which 


FIG. ii. 


comes  out  of  the  lungs  is  different  from  the  air  which  is  drawn  into 
them,  because  it  will  not  allow  a  candle  to  burn  in  it.1 

1  The  nature  of  this  gas  will  be  described  later  on. 


26 


Simple  Manipulations. 


Some  gases  cannot  be  collected  over  water,  because  either 
they  dissolve  in  water,  or  they  enter  into  chemical  combination 
with  water.  In  such  cases  the  mercurial  pneumatic  trough  may 
be  used  :  that  is  to  say,  the  liquid  metal  mercury  is  used  instead 
of  water,  provided  the  gas  is  without  chemical  action  upon 
mercury. 

When  a  gas  happens  to  be  either  much  lighter  or  much 
heavier  than  air,  we  may  collect  it  without  using  a  pneumatic 
trough  at  all. 

Experiment  32. — Obtain  a  ground  glass  plate  with  a  hole  in  the 
centre,1  and  fit  the  hole  with  a  cork  carrying  two  bent  tubes,  one 


FIG.  12. 


FIG.  13. 


FIG.  14. 


long,  and  the  other  quite  short,  as  shown  in  Fig.  14.  This  is 
supported  on  a  ring,  and  the  cylinder  to  be  filled  is  stood  over  the 
tubes.  Attach  the  long  tube  to  the  ordinary  coal-gas  supply, 
and  turn  on  the  gas.  Coal-gas  is  very  much  lighter  than  air  (the 
fact  that  it  is  used  for  filling  balloons  shows  this),  and  therefore,  by 
delivering  it  right  up  to  the  top  of  the  cylinder,  it  accumulates  in 
the  upper  part  of  the  vessel,  and  gradually  displaces  the  whole  of 
the  air,  driving  it  down  through  the  other  tube.  We  can  tell 

1  A  stout  piece  of  cardboard  will  answer  the  purpose,  although  the 
glass  plate  is  preferable. 


Collecting  Cases.  27 

when  the  air  is  all  displaced  by  smelling  the  gas  which  will  then  be 
escaping  at  the  exit  tube. 

This  method  of  collecting  is  called  upward  displacement. 

Gases  that  are  heavier  than  air  can  be  collected  by  down- 
ward displacement.  This  is  exactly  the  reverse  of  upward 
displacement,  and  is  carried  out  by  simply  inverting  the 
apparatus  in  Fig.  14,  so  that  the  cylinder  stands  mouth 
upwards. 


CHAPTER    IV. 

FITTING    UP    APPARATUS. 

MUCH  of  the  apparatus  required  for  such  experiments  as  the 
elementary  student  will  make,  he  can  himself  easily  put 
together  by  means  of  glass  and  rubber  tubes,  corks,  and  glass 
bottles.  Figs.  15,  16,  and  17  show  three  typical  forms  of 
apparatus  used  for  the  preparation  of  various  gases,  and  every 
student  should  fit  up  these  for  himself. 

Glass   tubes. — Glass    tubing  is   manufactured   in  many 


FIG.  15. 


varieties  of  glass,  and  of  many  sizes.  For  general  purposes  it 
is  best  to  obtain  soft  soda  glass ;  and  a  most  useful  size  is 
shown  in  section  at  0,  in  Fig.  18. 

Such  tube  is  cut  to  any  required  length  by  making  a  slight 
scratch  upon  it  with  the  edge  of  a  fine  three-cornered  file,  and 


Fitting  up  Apparatus. 


29 


then  breaking  it  across  exactly  as  one  would  snap  a  dry  twig 
of  wood. 

In  order  to  bend  a  tube,  it  is  made  soft  by  being  heated 
in  a  common  flat  gas-flame. 

Experiment  33. — Hold  a  short  piece  of  tube  in  the  flame  in  the 
manner  shown  in  Fig.  19,  and  slowly  rotate  it  between  the  thumbs 
and  forefingers  in  order  that  it  may  get  equally  heated  all  round. 


FIG.  16. 


FIG.  17. 


As  soon  as  it  is  quite  soft,  "withdraw  it  from  the  flame,  and  deliber- 
ately bend  it  to  a  right  angle  (Fig.  20).  When  cold,  wipe  off  the 
soot.  (Such  a  right  angle  bend  will  be  required  at  a,  a',  a"  Figs. 
1 6  and  17.) 

Tubes  should  not  be  heated  for  bending  in  a  Bunsen 
flame,  as  the  glass  in  this  case  will  become  creased  at  the  bend, 
as  in  Fig.  21,  B.  If  the 
tubing  employed  is  too  thin 
in  the  walls  (£,  Fig.  18),  the 
glass  will  collapse  at  the 
bend,  as  shown  in  Fig.  2 1, 
A.  Both  of  these  bends, 
besides  being  unsightly,  are  very  easily  broken,  and  they  also 
prevent  the  free  passage  of  gas  through  them. 

When  a  tube  is  to  be  bent  into  a  very  acute  angle,  it  should 
be  held  in  the  flame  in  the  manner  shown  in  Fig.   22,  with 


a 


FIG.  18. 


Fitting  up  Apparatus 


FIG  19 


FIG.  20. 


FIG  21. 


Bending  Glass  Tubes.  31 

the  knuckles  downwards  instead  of  upwards,  as  in  Fig.  19,  so 
that  the  bend  may  be  produced  by  bringing  the  hands  together 


FIG  22. 


FIG.  23. 


upwards,  as  in  Fig.   23.     Such  a  bend  is  required  in  making 
the  delivery  tub.es  d,  d,  d",  Figs.  16  and  17.     The  bent  tube 


32  Fitting  up  Apparatus. 

is  cut  at  r,  Fig.  24,  and  the  end  E  is  then  bent  as  shown  by 
the  dotted  lines,  by  holding  it  in  the  flame  as  explained  in 
Exp.  33.  All  glass  tubes,  when  bent  and  cut  to  shape 
and  length,  must  be  rounded  or  smoothed  at  their  extremities 
by  holding  them  in  a  flame  (Fig.  25)  until  the  "raw"  or  sharp 
edges  have  just  fused. 

"  Combustion "  tubing  is  made  of  hard  glass,  which  will 
stand  a  high  temperature  without  softening.     It  is  useful  when 
we  require    to    pass   a  gas 
over  some  heated   material. 
c.  Fig.  1 8,  shows  a  convenient 
size  for  combustion  tube. 

Corks. — Select  a  sound 
cork  which  is  just  too  large 
at  its  narrow  end  to  fit  into 
the  mouth  of  the  flask  or 


FIG.  24. 


FIG.  25. 


test- tube,  and  then  squeeze  it  either  in  a  cork- squeezer  or  by 
rolling  it  beneath  the  foot  on  the  floor,  using  a  moderate 
pressure  upon  it.  The  cork  should  then  fit  comfortably  into 
ihe  flask.  The  cork  has  next  to  be  bored  to  take  the  glass 
tube  which  is  to  pass  through  it.  This  is  best  done  by  means 
of  a  cork-borer,  a  tool  consisting  of  a  brass  tube,  with  a 
sharpened  edge  at  one  end ;  it  is  bored  into  the  cork 
much  in  the  same  way  as  a  gimlet  or  brad-awl  is  used  to 
bore  a  hole  in  wood,  except  that  the  cork-borer  is  usually 
wetted. 

In  order  to  be  sure  that  the  borer  selected  will  make  a 
hole  the  exact  size  required  for  the  glass  tube,  a  trial   hole 


Fitting  up  Apparatus.  33 

should  be  bored  in  a  waste  scrap  of  cork.  The  hole  must  be 
of  such  a  size  that  a  little  gentle  force  is  required  to  push  the 
tube  through  it.  When  using  the  cork-borer,  the  tool  should 
be  driven  in  at  the  narrow  end  of  the  cork,  care  being  taken 
to  make  the  hole  in  the  centre,  and  to  keep  it  straight.  If 
the  borer  is  fairly  sharp,  the  hole  will  be  clean  and  smooth. 

Sometimes  corks  are  bored  by  means  of  a  round  (or  rat- 
tail)  file;  but  this  is  not  a  good  plan,  as  it  is  much  more 
difficult  to  make  the  hole  perfectly  round ;  and  if  it  is  not 
round,  the  glass  tuj)e  will  not  fit  properly,  and  the  apparatus 
will  leak  at  this  point. 

If  the  bent  tube  has  been  rounded  at  its  ends  (p.  32),  it 
can  be  pushed  into  the  cork  without  cutting  or  tearing  the 
hole,  and  a  tight  fit  will  be  made. 

The  delivery  tube  is  attached  to  the  exit  tube  by  means  of 
a  short  piece  of  indiarubber  tube.  The  ends  of  both  glass 
tubes  being  smoothed,  as  already  described,  they  can  easily 
be  pushed  into  the  caoutchouc  connection  without  cutting  it, 
especially  if  the  latter  be  moistened  inside  by  breathing  through 
it  immediately  before  introducing  the  glass. 

The  apparatus,  type  No.  2  (Fig.  16),  differs  only  in  being 
fitted  with  a  cork  bored  with  two  holes,  into  the  second  of 
which  there  is  fitted  a  tube  known  as  a  thistle  Junnel.  This 
reaches  nearly  to  the  bottom  of  the  flask,  so  that  it  may  dip 
into  the  liquids  which  will  be  present ;  and  thus,  while  allowing 
liquids  to  be  poured  in,  will  prevent  gas  from  escaping  through 
the  funnel. 

The  pieces  of  apparatus  of  the  rirst  and  second  types  are 
intended  to  be  used  in  certain  experiments  when  heat  is 
applied,  hence  they  have  to  be  supported  by  suitable  clamps 
or  stands  at  a  convenient  height,  and  therefore  the  delivery 
tubes  must  be  made  of  such  a  length  that  they  will  reach  into 
the  pneumatic  trough. 

The  apparatus  of  type  No.  3  (Fig.  17)  is  used  in  the 
preparation  of  certain  gases,  when  it  is  not  necessary  to  apply 
any  heat  to  the  materials.  The  generating  vessel  in  this  case, 
instead  of  being  a  thin  glass  flask,  may  be  either  a  two-necked 
bottle,  B,  Fig.  17,  known  as  a  "  Woulfs  bottle,"  or  an  ordinary 

D 


34  Fitting  up  Apparatus. 

wide-mouthed  bottle.  In  the  latter  case,  a  single  cork  with  two 
holes  is  fitted  with  two  tubes  as  in  A,  Fig.  17  ;  but  when  the 
two-necked  bottle  is  used  the  thistle  funnel  is  fitted  into  one 
neck,  and  the  exit  tube  to  the  other.  This  form  of  apparatus 
is  preferable  to  the  other,  as  small  corks  are  more  easily  fitted 
so  as  to  be  free  from  leakage  than  wide  ones. 

When  the  apparatus  has  been  put  together,  and  before 
being  used,  it  should  be  tested  to  ascertain  whether  it  is  tight. 
In  the  case  of  A  and  B,  Fig.  15,  this  is  done  by  sucking  a 
little  of  the  air  out  of  the  apparatus  by  applying  the  mouth  to 
the  end  of  the  delivery  tube,  and  instantly  closing  the  tube 
with  the  tip  of  the  tongue.  If  the  joints  are  tight  the  tongue 
will  remain  drawn  to  the  end  of  the  tube,  and  a  little  effort  is 
felt  in  pulling  it  away;  whereas  if  there  is  any  leakage  in  the 
apparatus  the  tongue  parts  at  once  away  from  the  tube. 

To  test  the  other  forms  of  apparatus  in  Figs.  16  and  17, 
a  quantity  of  water  should  first  be  poured  into  the  flask  or 
bottle,  until  the  end  of  the  thistle  tube  dips  into  the  liquid. 
Then  by  applying  the  mouth  to  the  end  of  the  delivery  tube 
and  blowing  gently  into  the  apparatus,  the  water  will  be  forced 
up  the  funnel  tube.  It  should  in  this  way  be  driven  nearly  up 
to  the  head  of  the  thistle  tube,  and  the  end  of  the  delivery 
tube  closed  with  the  tongue.  If  the  apparatus  does  not  leak, 
the  water  will  remain  steady  in  the  thistle  tube,  otherwise  it 
will  gradually  sink  down.  If  the  apparatus  leaks  it  should  be 
refitted,  and  on  no  account  should  a  leaky  apparatus  be 
tight  by  the  use  of  scaling  wax  or  other  lutes. 


CHAPTER   V. 

SIMPLE   GLASS-BLOWING    OPERATIONS. 

GLASS-BLOWING  operations,  as  a  general  rule,  require  skill, 
practice,  and  patience  to  perform  with  anything  like  success ; 
nevertheless  there  are  a  number  of  smaller  and  simpler 
operations  which  can  easily  be  done  by  the  young  student, 
and  which  it  is  most  useful  that  he  should  know  how  to 
perform.  The  bending  of  glass  tube  and  the  rounding  of  the 
ends  have  already  been  described.  For  ihese  operations, 
however,  an  ordinary  gas-flame  is  employed,  but  for  those  now 
to  be  described  the  blow-pipe  is  to  be  used. 

A  blow-pipe  and  some  sort  of  foot-blower  usually  form  a 
part  of  the  regular  fittings  of  a  chemical  laboratory,  but  even 
in  their  absence  much  may  be  done  by  means  of  a  small 
Herapath  mouth  blow-pipe. 

General  rules. — (i)  Never  bring  a  piece  of  cold  glass 
directly  into  the  blow-pipe  flame,  but  first  warm  it  in  the  smoky 
flame  before  admitting  wind  from  the  blower. 

(2)  When  a  tube  is   being  heated,  it  should   (except   in 
special  cases)  be  continuously  revolved  in  the  flame,  so  that 
the  heating  may  be  uniform;    and  also,  as   it  gets   soft,  to 
prevent  the  glass  from  falling  out  of  shape. 

(3)  When  actually  blowing  glass,  always   remove  the  soft 
glass  from  the  flame. 

(4)  Always  begin  the  blowing  gently,  and  then  regulate  the 
force  of  the  breath  as  the  soft  glass  gives  to  the  pressure. 

To  open  out  the  end  of  a  glass  tube. — When  a  cork 
is  to  be  fitted  into  the  end  of  a  glass  tube,  the  end  should  be 
opened  out  a  little,  or  "  bordered."  The  simple  tool  required 


Simple  Glass-blowing  Operations. 


for  this  is  a  round  stick  of  charcoal,  pointed  at  one  end  like  a 
lead  pencil. 

Experiment  34. — Take  a  piece  of  moderately  wide  glass  tube 
and  warm  one  end  in  the  smoky 
flame  (Rule  i).  Then  admit  wind 
into  the  flame,  and  hold  the  tube  in 
the  position  shown  in  Fig.  26,  re- 
volving it  all  the  time  (Rule  2). 

When  the  end  of  the  glass  is 
sufficiently  soft,  remove  it  from  the 
flame,  and  push  the  pointed  piece 
of  charcoal  into  it,  giving  a  screw- 
ing motion  to  the  charcoal  (Fig.  27). 
When  the  tube  is  sufficiently  opened, 
hold  it  in  the  smoky  flame  again, 
gradually  turning  down  the  gas. 
This  anneals  the  glass,  that  is,  cools 
it  slowly,  and  makes  it  less  liable 
FlG-  26-  to  crack  afterwards. 


FIG.  27. 

To  draw  down  a  glass  tube  to  a  jet. 

Experiment  35. — Heat  a  piece  of  tube  in  the  ( blowpipe  flame, 


FIG.  28. 


holding  the  glass  as  in  Fig.  19.     As  soon  as  it  is  soft,  remove  it 


Joining  Glass  Tubes.  37 

from  the  flame,  and  pull  gently,  revolving  each  end  slowly  at  the 
same  time.  The  glass  then  assumes  the  form  shown  at  a,  Fig.  28, 
the  walls  of  the  tapering  and  narrow  parts  being  very  thin.  Now 
heat  another  piece,  keeping  it  longer  in  the  flame  ;  observe  that  the 
glass  gradually  thickens  and  the  walls  fall  together  as  the  mass 
gets  softer  and  softer  (<£,  Fig.  28).  Keep  the  tube  quickly  revolving, 
or  the  soft  part  will  drop.  Very  gently  draw  the  ends  apart,  still 
revolving,  and  the  tube  will  take  the  shape  seen  at  c,  Fig.  28,  where 
the  tapering  and  narrow  parts  are  thick  walled.  The  tube  may 
then  be  cut  at  any  desired  point  on  the  narrow  part  by  means  of  a 
file  scratch. 

To  seal  up  the  end  of  a  glass  tube. 

Experiment  36. — Draw  out  a  piece  of  glass  tube,  as  in  a,  Fig.  28. 
Cut  it  off  at  the  dotted  line  e,  and  heat  the  narrowed  end  of  d  in 
the  blow-pipe  flame  until  it  closes  up,  when  it  presents  the  appear- 
ance shown  at  f,  Fig.  29. 
Then  heat  the  somewhat 
thickened  end  in  the  blow- 
pipe, and,  when  just  soft,  blow 
gently  into  the  tube  (note 
Rules  3  and  4),  slowly  re- 
volving the  glass  at  the  time. 
It  should  then  appear  as  seen 
in  gj  Fig.  29.  If  too  much 

pressure  was  used  in  blowing,  the  end  will  be  expanded  into  a 
swelling.  It  may,  however,  be  reduced  by  revolving  it  in  the  blow- 
pipe flame,  when  the  sides  of  the  enlarged  part  will  fall  together 
again. 

To  join  two  tubes. 

Experiment  37. — Take  two  pieces  of  tube  of  the  same  diameter, 
close  one  end  of  one  with  a  cork,  and  heat  the  opposite  end  in  the 
blow-pipe.  At  the  same 
time  heat  one  end  of 
the  other  tube.    When 
the  ends  are  soft,  bring 
the  two  together  with 
a  little  pressure.     This 
causes  them  to  adhere, 

and,  at  the  same  time,  slightly  to  bulge  out  at  the  junction,  as  at 
a,  Fig.  30.  Then,  while  still  soft,  blow  gently  into  the  tube,  drawing 


38  Simple  Glass -blowing  Operations. 

it  out  slightly  at  the  same  time  so  as  to  keep  the  outer  walls 
parallel.  If  the  blowing  operation  has  not  followed  the  first  suffi- 
ciently quickly,  the  joint,  as  it  is  at  a,  may  be  re-heated  in  a  fine- 
pointed  blow-pipe  flame,  and  then  blown  as  described.  The  joint 
should  have  the  appearance  shown  at  b,  Fig.  30. 

Experiment  38. — Join  a  wide  tube  to  a  narrow  one.  First  draw 
out  the  wider  tube,  and  cut  it  off  when  it  has  a  diameter  equal  to 
the  narrow  tube.  Then  heat  the  drawn-out  end  of  the  wide  tube, 
and  one  end  of  the  narrow  tube,  and  join  them  in  the  way  de- 
scribed in  Exp.  37,  blowing  through  the  narrower  tube. 

To  seal  platinium  wire  into  glass  tubes,  (a)  In  the 
end  of  narrow  tubes. 

Experiment  39. — Draw  the  tube  out  to  a  point,  and  cut  it  oft 
so  that  the  wire  can  just  pass  through  the  drawn-out  end.  Then 

introduce  the  tip  into  a  flame,  when 

|p  ^K the  edges  of  the  glass  will  close  to- 

gether round  the  wire,  as  in  Fig.  31. 
Fio  ^  While  the  glass    is    still    soft,   the 

position  of  the  wire  can  be  adjusted, 
so  as  to  get  it  quite  straight. 

(b)  In  either  the  end,  or  the  side  of  a  wide  tube. 

Experiment  40. — Heat  the  glass,  at  the  point  where  the  wire  is 
to  be  introduced,  with  a  fine-pointed  flame,  and  when  a  small  spot 
is  soft,  stick  the  end  of  a  platinum  wire  into  it  and  draw  out  gently. 
In  this  way  a  tiny  branch  tube  is  made,  «,  Fig.  32.  This  is  then 
cut  off  short,  as  at  b.  The  wire  is  then  inserted,  and  the  flame 
again  directed  upon  it,  when,  as  in  Exp.  39,  the  glass  closes  up 
round  the  wire,  c,  Fig.  32. 

To  blow  a  bulb  on  the  end  of  a  tube. 

Experiment  41. — First  seal  up  the  tube  as  shown  at  /,  Fig.  29. 
Then  heat  the  extreme  end  where  the  glass  is  thick,  and  gently 
blow  it  out  so  as  to  obtain  the  result  seen  at  a,  Fig.  33.  Next  hold 
the  tube  in  a  large  blow-pipe  flame,  heating  mostly  the  part  between 
the  dotted  lines  ;  and,  as  the  glass  softens,  keep  quickly  revolving  it. 
It  then  assumes  somewhat  the  shape  seen  at  b,  Fig.  33.  Withdraw 
it  from  the  flame,  and  blow  steadily  into  the  tube,  holding  it  in  a 
horizontal  position,  and  revolving  the  glass  all  the  time,  ct  Fig.  33. 


Blowing  Bulbs. 


39 


[Probably  the  first  attempts  will  either  be  complete  failures,  or  very 
remarkably  shaped  bulbs,  but  with  a  little  patience,  better  results 
will  soon  follow.]  To  blow  a  larger  bulb,  more  glass  is  necessary, 


FIG.  32. 


FIG.  33. 


and  it  is  better  to  first  join  on  a  piece  of  larger  tube  as  described 
in  Exp.  38. 

To  blow  a  bulb  on  the  middle  of  a  tube. 

Experiment  42. — Close  one  end  of  the  tube  with  a  cork,  and  heat 
in  a  large  flame  at  the  spot  where  the  bulb  is  to  be  blown.  As  the 
glass  softens  and  thickens,  gently  press  it  together  so  as  to  ac- 


FlG.  34. 

cumulate  material  for   the  bulb,   Fig.   34.     Then   blow  steadily, 
holding  the  tube  in  a  horizontal  position,  and  revolve  it  between 

the  fingers. 


CHAPTER   VI. 

HYDROGEN. 

ALTHOUGH  solids  and  liquids  are  more  familiar  to  us  than 
gases,  it  will  nevertheless  be  more  convenient  to  begin  the 
study  of  chemical  facts  by  considering  some  of  the  methods 
of  preparing,  and  a  few  of  the  more  prominent  properties 
of  the  gaseous  element  hydrogen.  Let  us  understand  at  the 
very  outset,  that  chemists  cannot  create;  they  cannot  make 
hydrogen  from  either  nothing,  or  from  materials  which  have 
got  no  hydrogen  in  them.  The  alchemists  of  old  believed  in 
the  transmutation  of  the  metals,  and  they  spent  their  lives 
endeavouring  to  change  the  common  metals  into  gold ;  now- 
adays we  know  how  futile  the  attempts  were,  and  we  no  more 
expect  to  convert  copper  into  gold,  than  we  expect  to  "  gather 
figs  from  thistles." 

All  that  the  chemist  can  do  is  to  so  experiment  upon 
compounds  which  contain  hydrogen  as  one  of  their  constituents, 
as  to  cause  chemical  changes  to  take  place  which  will  result 
in  the  expulsion  of  this  hydrogen.  We  must,  therefore,  select 
some  suitable  compounds  containing  hydrogen  from  which  to 
obtain  this  element. 

Chemists  have  found  out,  by  numberless  experiments,  that 
hydrogen  forms  one  of  the  constituents  of  a  vast  number  of 
compounds;  thus  it  is  found  to  be  present  in  nearly  all 
animal  and  vegetable  substances.  It  is  a  constituent  of  water, 
and  also  of  all  those  things  chemists  term  acids.  There  are 
three  common  compounds  from  which  the  element  is  most 
usually  obtained ;  these  are— (i)  Water,  (2)  Sulphuric  Acid, 
(3)  Hydrochloric  Acid. 


Hydrogen.  4 1 

(i)  Hydrogen  from  Water. — As  water  is  the  commonest 
of  these  three  substances,  we  will  first  experiment  with  it. 
Water  is  composed  of  the  two  elements,  hydrogen  and  oxygen, 
chemically  united  together,  and  in  order  to  separate  them  by 
chemical  processes,  we  must  find  some  element  which,  under 
suitable  conditions,  can  overcome  the  force  uniting  the  oxygen 
and  hydrogen,  some  element  which  can  seize  the  oxygen  and 
tear  it  away,  so  to  say,  from  the  grasp  of  the  hydrogen.  Many 
metals  are  capable  of  doing  this. 

Experiment  43.— Carefully  throw  a  small  fragment  of  sodium, 
about  the  size  of  a  pea,  upon  some  cold  water  in  an  ordinary 
dinner-plate.  Notice  that  the  metal  at  once  melts,  and  the  globule 
swims  to  and  fro  upon  the  surface  of  the  water  (much  in  the  same 
manner  that  a  drop  of  water  runs  about  upon  a  hot  iron),  producing 
a  hissing  sound.  Observe  that  the  globule  quickly  gets  less  and 
less,  and  finally  disappears.  Now  take  a  strip  of  turmeric  paper 
(that  is,  blotting-paper  which  has  been  dyed  yellow  with  turmeric) 
and  dip  it  first  into  clean  water  ;  this  causes  no  stain  upon  it ;  now 
dip  it  into  the  water  in  the  plate,  and  observe  that  it  is  strongly 
stained  brown.  Also  dip  the  fingers  into  the  water,  and  notice  that 
it  feels  slimy,  or  caustic,  to  the  touch.  This  shows  that  there  is 
something  in  the  water  after  the  sodium  has  been  in  contact  with 
it  which  was  not  previously  there.  One  of  the  products,  there- 
fore, of  the  action  of  the  metal  sodium  upon  water  is  something 
which  dissolves  in  the  excess  of  water  present,  yielding  a  solution 
which  turns  yellow  turmeric  brown,  and  is  caustic  to  the  touch. 
The  other  product  of  the  action  was  hydrogen  gas,  which  escaped 
unnoticed  into  the  air.  If  we  repeat  the 
experiment,  using  potassium  instead  of  so- 
dium, the  hydrogen  will  not  escape  our 
observation. 

Experiment  44. — Throw  a  similar  frag- 
ment of  potassium  upon  some  clean  water 
in  a  plate,  and  at  once  cover  the  whole  with 
a  glass  bell  jar,  as  in  Fig.  35.  The  potassium 
appears  to  take  fire  the  moment  it  touches  FlG 

the  water,  but  really  it  is  the  hydrogen  which 

burns,1  the  hydrogen  which  is  driven  out  of  its  combination  with 
oxygen.      The  action  of  the  potassium  upon  water  develops  so 

1  A  little  of  the  potassium  burns  also,  and  this  gives  to  the  flame  of  the 
burning  hydrogen  the  violet  colour* 


42  Hydrogen. 

much  heat  as  to  set  fire  to  the  hydrogen.  Carefully  notice  that  for 
a  few  seconds  after  the  flame  goes  out,  a  little  red-hot  globule  of 
something  in  a  melted  state  remains  swimming  upon  the  water,  and 
then  suddenly  disappears  with  a  little  splutter.  (//  is  in  order  to 
Prevent  this  substance  from  being  scattered,  and  injuring  the  eyes, 
that  the  plate  must  be  covered  with  the  bell  glass?)  Test  the  water 
in  the  plate  with  turmeric  paper,  and  the  same  stain  is  produced  as 
in  the  former  experiment.  Therefore,  when  potassium  acts  upon 
water,  hydrogen  gas  is  expelled,  and  a  substance  is  formed  which 
also  stains  turmeric. 

We  must  now  adopt  some  device  for  collecting  fat  hydrogen 
gas  which  can  thus  be  expelled  from  its  combination  with 
oxygen,  by  the  action  of  sodium  upon  the  water.  There  are 
various  ways  of  doing  this,  but  only  one  of  them  is  free  from 
danger ;  for  sodium,  when  brought  carelessly  into  contact  with 
water,  is  liable  to  give  rise  to  a  serious  explosion. 

Experiment  45. — Take  a  piece  of  lead  pipe,  2\  centimetres  long 
(i  inch)  and  £  centimetre  bore,  and  close  up  one  end  by  hammering 
the  lead,  as  seen  in  Fig.  36.  This  little  pipe  is  then  filled 
with  sodium  by  first  rolling  a  pellet  of  the  metal  (which 
is  about  as  soft  as  wax)  between  the  fingers  until  it 
will  just  push  into  the  tube,  and  then  forcing  it  in  by 
pressing  the  mouth  of  the  tube  firmly  down  upon  the 
table.  The  excess  of  sodium  round  the  mouth  is  then 
trimmed  off  with  a  knife.  When  this  tube,  with  its  little 
G<  3  '  charge  of  sodium,  is  dropped  into  water  in  a  dish  or 
trough,  it,  of  course,  sinks  to  the  bottom,  and  a  stream  of  gas- 
bubbles  will  be  seen  rising  through  the  water  ;  and  if  the  mouth 
of  an  inverted  glass  cylinder,  filled  with  water,  be  brought  over  the 
ascending  bubbles,  as  in  Fig.  37,  the  gas  will  collect  in  the  upper 
part.  If  the  sodium  is  all  used  up  before  the  cylinder  is  full,  a 
second  leaden  tube  can  be  placed  in  the  water.  Now  remove  the 
cylinder,  as  described  (page  25),  and  apply  a  lighted  taper  to  the 
mouth  of  the  jar.  Notice  that  the  gas  burns  quickly,  with  rather 
a  yellowish  flame,  this  yellow  colour  being  due  to  a  trace  of 
sodium. 

If  the  water  in  the  trough  be  tested  with  turmeric  paper, 
it  will  stain  it  brown,  as  in  the  former  case.  This  other 
product  of  the  action  of  sodium  on  water,  which  remains 
dissolved  in  the  large  excess  of  water  in  the  dish,  is  called 


Hydrogen  from    Water.  43 

sodium   hydroxide,   or   caustic  soda ;    while  that  which  was 
formed  in  the  case  of  potassium  was  potassium  hydroxide. 

Many  other  metals  besides  sodium  and  potassium  are 
able  to  expel  the  hydrogen  from  water  if  the  conditions  are 
favourable.  Thus — 

Experiment  46. — Drop  a  few  fragments  of  magnesium  into  cold 
water  in  a  test  tube,  observe  that  no  action  takes  place  between 


FIG.  37. 

the  metal  and  the  water  ;  in  this  respect,  therefore,  magnesium  is  a 
contrast  to  sodium  or  potassium.  Now  boil  the  water,  and  notice  that 
even  at  the  boiling  temperature  practically  no  hydrogen  is  given  off. 

This  experiment  shows  that  magnesium  is  not  able  to 
decompose  water  appreciably  even  at  the  boiling  point.  Let 
us,  therefore,  try  the  effect  of  heating  the  metal  much  more 
strongly  in  a  current  of  steam  :  that  is,  of  heating  it  in  contact 
with  water  in  the  state  of  gas. 

Experiment  47. — Fold  up  a  short  strip  of  magnesium  ribbon, 
and  place  it  in  a  hard  glass  bulb  which  is  held  in  a  slightly 
inclined  position  by  a  clamp,  as  in  Fig.  38.  Attach  to  one  end  of 


44 


Hydrogen. 


the  bulb  tube  a  small  empty  flask,  in  the  manner  shown  in  the 
figure,  and  connect  this  to  a  tin  can  in  which  water  is  being  boiled. 
As  soon  as  steam  issues  from  the  bulb  tube,  gradually  heat  the 
latter  all  along,  by  means  of  a  Bunsen  flame,  until  the  inside  is 
quite  dry,  and  the  steam  no  longer  condenses  as  it  passes  through. 
Then  hold  the  flame  steadily  under  the  magnesium,  and  heat  it 
until  it  is  nearly  red  hot,  when  it  will  suddenly  take  fire  and  burn 
in  the  steam,  combining  with  the  oxygen  of  the  steam  and  letting 
the  hydrogen  go  free.  Directly  this  takes  place,  light  the  hydrogen 
as  it  escapes  from  the  end  of  the  tube.  Notice  that  it  burns 
with  a  scarcely  visible  flame,  because  magnesium  does  not  impart 


FIG.  38. 

any  colour  to  a  hydrogen  flame  like  sodium  and  potassium  do. 
Notice  also  that  the  residue  in  the  bulb  is  white. 

Shake  out  some  of  this  and  see  if  it  will  dissolve  in  water  by 
adding  water  to  a  little  in  a  test-tube.  Note  that  it  is  not  percep- 
tibly soluble.  Take  a  little  more  of  it  and  place  it  on  a  piece  of 
turmeric  paper,  and  moisten  it  with  a  drop  of  water ;  notice  that 
the  paper  is  stained  slightly  brown,  showing  a  slight  resemblance 
between  the  behaviour  of  this  substance  (magnesium  oxide  or 
magnesia)  and  sodium  hydroxide. 

The  common  metal  iron,  at  a  bright-red  heat,  will  also 
expel  the  hydrogen  from  water  when  the  latter  is  in  the 
condition  of  water  vapour  or  steam.  This  method  of  obtain- 
ing hydrogen  is  often  made  use  of  on  a  large  scale. 

Experiment  48. — Take  a  piece  of  ordinary  iron  gas-pipe,  suffi- 
ciently long  to  project  15  cm.  (6  inches)  beyond  each  end  of  the 
furnace  to  be  used.  Fill  the  pipe  with  small  iron  nails,  and  fit 
each  end  with  a  cork  with  a  short  straight  glass  tube.  To  one 


Hydrogen  from  Sulphuric  Acid.  45 

end  a  delivery  tube  is  attached,  while  the  other  is  connected  to  a 
small  steam  boiler  (conveniently  made  out  of  a  common  tin  can, 
see  Fig.  38).  The  iron  pipe  is  heated  to  a  bright-red  heat,  either 
in  a  gas  or  coke  furnace,  and  steam  is  passed  through  it.  The  red 
hot  iron  seizes  the  oxygen  of  the  steam,  combining  with  it  to  form 
iron  oxide,  which  remains  in  the  tube,  while  a  rapid  stream  of 
hydrogen  is  evolved,  and  this  may  be  collected  over  water  in  the 
pneumatic  trough. 

(2)  Hydrogen  from  Sulphuric  Acid. — Although  water 
is  the  commonest  compound  of  hydrogen,  sulphuric  acid  is 
the  one  from  which  it  is  most  convenient  to  obtain  hydrogen, 
and  when  we  require  this  gas  for  experiments,  it  is  almost 
always  got  from  sulphuric  acid  by  the  action  upon  it  of  the 
metal  zinc, 

Experiment  49. — Place  some  granulated  zinc  x  in  a  two-necked 


FIG.  39. 

Woulf's  bottle,  arranged  as  in  Fig.  39,  with  just  sufficient  water 
to  cover  it,  and  pour  upon  it  by  means  of  the  thistle  funnel  a  little 
strong  sulphuric  acid.  Almost  immediately  it  will  be  noticed  that 
an  effervescence  begins,  showing  that  gas  is  being  disengaged. 
Allow  the  action  to  go  on  for  a  few  minutes  until  all  the  air  origi- 
nally present  in  the  bottle  has  been  swept  out  by  the  gas,  and  then 
collect  the  gas  at  the  pneumatic  trough.  Three  or  four  cylinders 
should  be  filled  for  subsequent  experiments. 

In  this  experiment   the   metal  zinc  expels  the  hydrogen 
from  sulphuric  acid,  and  combines  with  what  is  left  of  the 

1  Granulated  zinc  is  made  by  first  melting  the  metal  and  pouring  it 
in  a  thin  stream  into  cold  water. 


46  Hydrogen. 

sulphuric  acid  after  the  hydrogen  is  gone.  The  compound 
so  produced,  called  zinc  sulphate  (or  white  vitriol),  is  left 
dissolved  in  the  water  in  the  bottle.  It  may  be  obtained 
from  the  solution  by  the  following  experiment. 

Experiment  50. — When  the  above  experiment  is  concluded, 
pour  the  liquid  out  of  the  Woulf 's  bottle  and  filter  it.  Then  gently 
evaporate  it  in  a  dish  over  a  small  flame  until  it  is  reduced  to 
about  half  the  bulk,  and  allow  it  to  cool,  when  long  glassy-like 
colourless  crystals  will  deposit.  This  is  the  zinc  sulphate. 

Instead  of  employing  zinc  in  Exp.  49,  the  metal  iron 
might  have  been  used ;  but  the  hydrogen  obtained  in  this 
way  is  always  contaminated  with  compounds  of  hydrogen 
with  both  carbon  and  sulphur  (both  of  these  elements  always 
being  present  in  ordinary  iron),  which  give  to  the  hydrogen 
an  unpleasant  smell. 

Experirnent  51. — Put  a  small  quantity  of  iron  filings  into  a  test- 
tube,  and  pour  upon  them  a  little  dilute  sulphuric  acid.  Notice 
effervescence,  due  to  the  escape  of  gas.  Smell  the  gas,  and  note 
the  peculiar  and  nasty  odour.  This  smell  does  not  belong  to 
hydrogen,  but  to  the  impurities  present.  Bring  a  lighted  taper 
to  the  mouth  of  the  tube  ;  observe  that  the  gas  burns.  When  the 
action  has  continued  for  some  time,  the  liquid  may  be  filtered,  and 
the  clear  solution  evaporated,  when  small  green  crystals  will  be 
deposited.  This  substance  is  the  compound  of  iron  with  what  is 
left  of  the  sulphuric  acid  after  the  hydrogen  in  it  is  expelled. 
It  is  called  ferrous  sulphate  (or  green  vitriol). 

In  a  similar  manner  we  can  expel  the  hydrogen  from 
sulphuric  acid  by  means  of  the  metal  magnesium. 

Experiment  52. — Drop  a  few  fragments  of  magnesium  ribbon 
into  a  little  dilute  sulphuric  acid  in  a  test-tube.  Notice  how  briskly 
the  gas  is  given  off.  Show  that  it  is  hydrogen  by  lighting  it  at  the 
mouth  of  the  tube.  When  the  metal  is  all  dissolved,  this  solution 
also  may  be  evaporated  down  and  allowed  to  cool,  when  colourless 
crystals  of  magnesium  sulphate  will  be  formed.  Compare  these 
crystals  with  those  of  zinc  sulphate. 

(3)  Hydrogen  from  Hydrochloric  Acid.— The  hydrogen 
from  this  compound  is  also  expelled  by  the  metals,  zinc, 
iron,  and  magnesium ;  therefore,  hydrochloric  acid  may  be 


The  Properties  of  Hydrogen.  47 

used  instead  of  sulphuric  acid  in  Exps.  49  to  52.  The 
compounds  which  remain  behind  in  the  solution  would,  in 
this  case,  be  compounds  of  the  metals  with  what  is  left  of 
hydrochloric  acid  after  the  hydrogen  has  been  expelled ;  they 
would  be  zinc  chloride,  ferrous  chloride,  and  magnesium 
chloride  respectively. 

The  Properties  of  Hydrogen.— From  the  various 
samples  collected,  we  learn  that  the  gas  is  colourless  (there- 
fore invisible).  Also  that  it  does  not  appreciably  dissolve 
in  water,  for  when  left  standing  in  vessels  in  the  pneumatic 
trough  the  water  does  not  show  any  signs  of  rising  up  in  the 
cylinders,  which  it  would  if  the  gas  were  soluble.  If  we  take 
one  of  the  jars  of  gas  and  smell  it,  we  shall  find  that  it  is 
quite  odourless.  We  have  also  learnt  from  Exps.  45  and 
47  that  hydrogen  will  burn. 

It  is  important  to  remember  that  if  we  take  any  gas  which 
will  burn  in  the  air,  and  previously  mix  it  with  a  certain 
proportion  of  air,  and  then  bring  a  light  to  the  mixture,  an 
explosion  results.  Every  one  knows  that  it  is  dangerous  to 
bring  a  light  into  a  room  where  there  is  an  escape  of  coal-gas, 
that  is,  where  there  is  a  mixture  of  gas  and  air.  Coal-gas  alone 
burns  quietly ;  but  a  mixture  of  coal-gas  and  air  in  certain 
proportions  explodes  when  lighted.  This  is  particularly  true 
of  hydrogen.  We  have  seen  by  Exp.  47  that,  when  un- 
mixed with  air,  it  burns  quietly,  but  if  a  mixture  of  hydrogen 
and  air  be  lighted,  the  mixture  explodes  with  great  violence. 
Therefore,  before  bringing  a  flame  to  the  tube  leading  from 
a  hydrogen  apparatus,  it  is  most  important  to  be  sure  that  all 
the  air  originally  present  in  the  bottle  has  been  swept  out. 
In  order  to  realize  the  force  of  the  explosion  that  would  occur 
under  such  circumstances,  and  the  danger  which  would  result 
from  neglecting  this  precaution,  the  following  experiment  may 
be  made — 

Experiment  53. — Fit  an  ordinary  pear-shaped  soda-water  bottle 
with  a  cork,  through  which  is  fixed  a  short  piece  of  the  stem  of  a 
clay  tobacco-pipe,  only  just  projecting  through  the  cork.  Put  a 
little  granulated  zinc  into  the  bottle,  add  some  dilute  sulphuric 
acid,  and  insert  the  cork.  Hold  a  lighted  taper  to  the  end  of  the 


Hydrogen. 


tube  as  though  attempting  to  inflame  the  hydrogen,  and  in  a  few 
seconds,  when  the  gas  has  mixed  with  the  air  in  the  bottle  in  a 
certain  proportion,  a  loud  explosion  will  result,  which  will  shoot 
the  cork  out  of  the  bottle  with  some  violence.  There  is  no  fear 
of  the  thick  bottle  bursting,  but  it  should  be  held  in  such  a  position 
that  the  cork  will  not  fly  in  a  direction  where  it  can  do  harm. 

Experiment  54. — Disconnect  the  bent  delivery  tube  from  the 
hydrogen  apparatus,  and  by  means  of  a  rubber  tube  attach  a 
straight  glass  tube.  Fill  a  test-tube  with  the  gas  by  upward 
displacement,  as  in  Fig.  40.  After  a  few  moments  withdraw  the 

delivery  tube,  and  ap- 
ply a  light  to  the 
mouth  of  the  test- 
tube.  If  the  hydrogen 
is  mixed  with  air,  a 
slight  pop  will  result, 
but  if  free  from  air  it 
will  burn  quietly.  If 
this  is  the  case  the 
gas  may  be  lighted  at 
the  end  of  the  de- 
livery tube.  Notice 
that  the  flame  has  a 
yellowish  colour,  this 
is  due  to  the  soda 
present  in  the  glass 
(recall  the  appearance 
of  the  flame  when  the 
hydrogen  collected  in 
Exp.  45  was  burnt). 
The  true  appearance 
of  a  flame  of  hydrogen 
may  be  seen  by  sub- 
stituting for  the  glass  tube  a  short  piece  of  lead  pipe,  into  one  end 
of  which  an  ordinary  metal  gas-burner  has  been  screwed.  Notice 
that  the  flame  now  is  almost  without  colour,  being  slightly  bluish, 
and  gives  no  light. 

Place  an  ordinary  coal-gas  flame  by  the  side  of  the  hydrogen 
flame  and  compare  them.  Depress  into  each  flame  a  white  plate, 
observe  that  the  bright  flame  blackens  the  plate,  while  no  soot  is 
deposited  from  the  other.  Hold  a  clean  dry  jar  over  each  flame 
for  a  few  moments,  notice  that  in  each  case  the  flames  are  giving 


FIG.  40. 


The  Burning  of  Hydrogen.  49 

off  steam,  because  moisture  will  be  deposited  on  the  cold  sides  of 
the  glass.  Cover  each  of  the  jars  with  a  glass  plate,  or  piece  of 
card,  and  pour  into  each  a  little  clear  lime-water.  After  shaking 
the  lime-water  with  the  air  in  the  jars,  it  will  be  seen  that  in  the 
jar  which  was  held  over  the  gas  flame,  the  lime-water  has  turned 
milky,  while  in  that  which  was  over  the  hydrogen  it  remains 
clear. 

These  experiments  teach  us  that  water  is  formed  when  we 
burn  either  hydrogen  or  coal-gas  ;  and  that  from  burning  coal- 
gas  we  also  get,  besides  the  water,  a  gas  which  will  turn  lime-water 
milky. 

Hydrogen  is  the  lightest  substance  known  to  chemists. 
It  is  nearly  14^  times  lighter  than  air.  On  account  of  its 
extreme  lightness  it  can 
be  collected  in  vessels  by 
upward  displacement  (see 
Fig.  41).  It  may  also  be 
poured  upwards  from  one 
vessel  to  another. 

Experiment  55. — Take 
a  cylinder  containing  only 
air,  and  hold  it  mouth  down- 
wards   in   the    left    hand  ;  FIG.  41. 
then  take  another  cylinder 

filled  with  hydrogen  in  the  other  hand,  and  bring  its  mouth  just 
beneath  that  of  the  first,  in  the  manner  shown  in  Fig.  41,  and 
gradually  empty  its  contents  up  into  this  one  by  lowering  the  foot 
until  the  position  shown  in  Fig.  42  is  reached.  Now  stand  the 
lower  cylinder  down,  and  apply  a  light  to  the  contents  of  first 
one  and  then  the  other.  Notice  that  no  hydrogen  is  left  in  the  one 
which  formerly  was  full,  and  that  the  contents  of  the  upper  jar  burn 
with  the  characteristic  flame. 

Owing  to  the  extreme  lightness  of  hydrogen,  it  is  some- 
times used  for  filling  balloons.  On  a  small  scale  we  can 
imitate  this  by  filling  soap  bubbles  with  the  gas. 

Experiment  56. — Attach  a  common  clay  tobacco-pipe  to  the 
hydrogen  apparatus,  and  proceed  to  blow  a  bubble,  holding  the 
pipe  in  the  usual  position.  Notice  that  the  bubble  in  trying  to 

E 


5O  Hydrogen. 

ascend,  very  soon  begins  to  curl  over  on  to  the  outside  of  the  pipe, 
as  in  Fig.  43.     Turn  the  pipe  upside  down,  and  blow  another,  and 

observe  that  as  it  grows  it  struggles 
to  tear  itself  away  from  the  pipe, 
assuming  the  shape  seen  in  Fig.  44. 
When  disengaged  from  the  pipe  it 
ascends  very  quickly. 

A  pretty  experiment  is  to  make 
a  small  hydrogen  bubble  carry  up  a 
large  air  bubble.  To  do  this,  first 
blow  an  ordinary  bubble  with  the 
breath.  Then,  while  it  is  still  hang- 
ing on  the  pipe,  bring  under  it  a 
hydrogen  bubble  just  beginning  to 
form  (Fig.  45).  As  the  latter  increases 
in  size,  gradually  invert  the  two  pipes 
so  that  the  one  holding  the  hydrogen 
bubble  is  uppermost.  Then,  with  a 
slight  jerk,  detach  first  the  pipe  from 
the  air  bubble,  and  next  the  one  con- 
veying hydrogen.  The  double  bubble  then  slowly  rises  (Fig.  46), 
and  generally  when  it  touches  the  ceiling  the  little  hydrogen  one 
breaks,  and  allows  the  other  to  fall  to  the  ground  again.  This 
experiment  requires  a  good  soap  solution.1 


FIG.  42. 


FIG.  43. 

Although  hydrogen  will  burn,  it  will  extinguish  the  flame 
of  any  ordinary  burning  substance.  This  is  true  of  all  gases 

1  The  following  is  a  good  receipt.  Dissolve  2  grams  of  sodium  oleate 
in  80  cc.  cold  distilled  water,  add  20  cc.  glycerine,  well  shake,  and  put 
away  in  a  dark  cupboard  for  two  days  to  settle.  Then  carefully  pour 
off  the  clear  liquid,  add  just  I  drop  of  ammonia  and  shake  up. 


The  Lightness  of  Hydrogen. 


52  Hydrogen. 

which  burn  in  the  air,  that  they  will  put  out  the  flames  of 
all  other  things  which  bum  in  the  air. 

Experiment  57. — Thrust  a  lighted  candle  (or  piece  of  thick 
taper),  fastened  to  the  end  of  a  wire,  into  a  cylinder  of  hydrogen 
held    mouth    downward.      As   the  burning 
candle  approaches  the  mouth  of  the  vessel, 
it  there  sets  fire  to  the  hydrogen  ;  but  as  it 
is  pushed  up  into  the  gas  its  own  flame  is 
extinguished.      Withdraw    the    candle    and 
,.v  _  .      relight  it  as  it  passes  out  through  the  still 
/    *     .  burning  hydrogen. 

The  same  will  happen  if  a  paper  spill 
be  used   instead   of   a   taper,    or  if   an 
ordinary   coal-gas    flame   be    introduced. 
We    therefore    say   that    hydrogen    is  a 
combustible  gas,  but  that  it  is  a  non-supporter  of  combustion. 

EPITOME. 

The  element  hydrogen  is  found  free  in  only  very  small  quantities 
on  the  earth  ;  but  it  is  present  on  the  sun  in  enormous  quantities. 
It  occurs  in  nature  in  combination  with  oxygen  in  water  ;  with 
carbon  in  marsh  gas  ;  with  sulphur  in  sulphuretted  hydrogen.  It 
is  present  in  nearly  all  animal  and  vegetable  substances,  and  is  a 
constituent  of  all  acids. 

It  can  be  obtained  from  water  :  (i)  by  the  action  of  either 
sodium  or  potassium  at  the  ordinary  temperature  ;  (2)  by  the  action 
of  magnesium  or  iron  at  a  red  heat. 

It  may  be  got  from  either  sulphuric  acid  or  from  hydrochloric 
acid  by  the  action  of  either  zinc,  iron,  or  magnesium  upon  them. 

The  common  laboratory  method  is  by  acting  on  sulphuric  acid 
with  zinc. 

Hydrogen  is  a  colourless,  tasteless,  odourless  gas.  It  burns 
with  a  nearly  invisible  bluish  flame. 

A  mixture  of  hydrogen  and  air  explodes  when  lighted.  It 
extinguishes  the  flames  of  other  ordinary  burning  substances. 

Hydrogen  is  the  lightest  of  all  known  substances,  being  14-4 
times  lighter  than  air.  For  this  reason  it  is  taken  as  the  standard 
or  unit  for  comparing  the  densities  of  all  other  gases.  Thus  we 
say  that  the  density  of  air  is  14*4,  which  means  that  air  is  14-4 


Hydrogen.  53 

times  heavier  or  denser  than  hydrogen,  bulk  for  bulk.     Hydrogen 
being  thus  taken  as  the  unit,  obviously  its  density  is  unity  or  i. 
Reactions  for  hydrogen —  * 

(1)  From  water,  by  the  action  of  sodium  H2O  +  Na  =  NaHO  +  11. 

(2)  ,,  ,,  ,,         magnesium  H2O  +  Mg  =  MgO  +  H2. 

(3)  „  *,  „         iron4H2O  +  3Fe  =  Fe3O4  +  4H2. 

(4)  From  sulphuric  acid,  by  the  action  of  zinc  H2SO4  +  Zn  =  ZnSO.,  +  H2. 

(5)  „  ,,  ,,        magnesium  H2SO4  +  Mg  =  MgSO4  +  H2. 

(6)  ,,  ,,  ,,        iron  H2SO4  +  Fe  =  FeSO4  +  H2. 

(7)  From  hydrochloric  acid,  by  the  action  of  zinc  2  HC1  +  Zn  =  ZnCJ2  +  H2. 

(8)  The  combustion  of  hydrogen  in  air  or  oxygen  H2  +  O  =  H2O. 


1  These  signs  and  symbols  will  be  explained  later  on  ;  the  student  may 
pass  them  over  at  this  stage. 


CHAPTER   VII. 

OXYGEN. 

THIS  element,  like  hydrogen,  is  a  gas,  but  in  almost  every 
other  respect  it  presents  a  complete  contrast  to  that  element. 
There  are  a  great  many  compounds  containing  oxygen  as 
one  of  their  constituents,  from  which  the  element  can 
easily  be  obtained,  for  in  combination  with  others  it  is  the 
most  abundant  of  all  the  elements.  Unlike  hydrogen,  it 
is  found  in  the  free  or  uncombined  state  in  large  quan- 
tities on  the  earth,  for  the  atmosphere  consists  essentially 
of  free  oxygen,  mixed  with  about  four  times  its  volume  of 
nitrogen. 

We  shall  consider  the  methods  of  obtaining  this  element 
from  four  of  its  compounds,  namely,  from — 

(i)  Mercuric  oxide ;  (2)  Potassium  Chlorate;  (3)  Sodium 
peroxide ;  (4)  Water;  and  also  the  way  by  which  it  is  obtained 
from  the  air. 

(i)  Oxygen  from  Mercuric  Oxide.— We  saw  in  Exp.  6 
that  when  this  compound  is  simply  heated,  it  is  decom- 
posed into  its  two  constituent  elements,  oxygen  and  mercury. 
In  order  to  collect  the  gas  which  is  given  off  we  proceed  as 
follows — 

Experiment  58. — Heat  a  small  quantity  of  the  red  powder  in 
a  hard  glass  tube,  arranged  as  shown  in  Fig.  15,  A,  and  collect 
the  gas  over  water.  Notice  that  the  evolution  of  gas  is  not  very 
rapid  ;  also  that,  as  in  Exp.  6,  the  metal  mercury  collects  on  the 
cooler  part  of  the  tube  in  small  globules. 

The  chief  interest  in  this  experiment  lies  in  the  fact  that 


Oxygen.  55 

it  was  the  very  method  by  which  Priestley  first  discovered 
oxygen  in  1774.  He  called  it  dephlogistigated  air;  the  name 
oxygen  was  given  to  it  later  by  Lavoisier. 

(2)  Oxygen  from  Potassium  Chlorate. — When  this 
salt  is  heated,  it  first  melts  and  then  rapidly  gives  off  oxygen. 
It  is  decomposed  by  heat  into  two  things,  namely,  into  oxygen, 
which  passes  off  as  gas,  and  into  potassium  chloride,  which 
remains  behind  as  a  white  solid. 

Experiment  59. — Heat  a  small  quantity  of  potassium  chlorate 
in  a  similar  apparatus  to  that  used  for  the  last  experiment.  Notice 


FIG.  47. 

that  the  crystals  first  crackle,  then  melt,  and  that  the  melted 
compound  then  begins  to  effervesce,  owing  to  the  escape  of  the 
oxygen  gas.  Note  also  that  the  gas  is  given  off  much  more  rapidly 
than  in  the  case  of  the  mercuric  oxide. 

It  has  been  found  that  if  the  potassium  chlorate  be 
previously  mixed  with  manganese  dioxide,  the  chlorate  gives 
up  its  oxygen  much  more  rapidly,  and  at  a  much  lower 
temperature. 

Experiment  60. — Mix  about  20  grams  of  potassium  chlorate 
with  about  a  quarter  of  its  weight  of  powdered  manganese  dioxide, 
and  gently  heat  the  mixture  in  a  flask  (a  common  Florence  oil 
flask)  as  shown  in  Fig.  47.  Notice  first  that  the  mixture  does  not 


56  Oxygen. 

melt  ;  also  that  moisture  collects  in  the  neck  of  the  flask.  This 
chiefly  comes  from  the  manganese  dioxide,  which  is  always  damp  ; 
and  it  is  in  order  to  prevent  this  condensed  moisture  from  running 
back  into  the  heated  flask  and  cracking  it,  that  the  apparatus  is 
supported  in  a  horizontal  position.  Observe  how  rapidly  the  gas  is 
evolved,  and  with  what  a  little  heat.  It  is  to  allow  plenty  of  passage 
for  the  gas  that  the  usual  glass  delivery  tube  is  here  replaced  by  a 
wide  piece  of  indiarubber  tube. 

Notice  that  during  the  experiment  little  sparkles  occasionally 
appear  in  the  heated  mixture.  These  are  caused  by  small  particles 
of  combustible  impurities  which  are  always  liable  to  be  present  in 
manganese  dioxide.1 

Collect  several  cylinders  or  jars  with  the  gas  as  it  is  evolved, 
and  keep  them  standing  mouth  downward  in  small  plates  containing 
a  little  water. 

This  is  the  method  usually  employed  in  the  laboratory  for 
preparing  oxygen. 

The  material  left  in  the  flask  after  the  experiment  consists 
of  potassium  chloride  and  manganese  dioxide.  The  latter 
substance  therefore  comes  out  of  the  reaction  in  exactly  the 
same  slate  as  it  was  at  the  beginning.  If  the  mixture  is  boiled 
with  water,  the  potassium  chloride  dissolves  and  leaves  the 
manganese  dioxide ;  so  that  if  the  mixture  is  then  filtered,  the 
dioxide  can  be  recovered,  and  used  over  and  over  any  number 
of  times.  The  way  in  which  it  acts  in  causing  the  potassium 
chlorate  to  give  up  its  oxygen  more  readily,  involves  a  series 
of  rather  complex  changes,  which  cannot  be  conveniently 
considered  at  this  stage. 

(3)  Oxygen  from  Sodium  Peroxide. — When  this  com- 
pound is  brought  into  contact  with  water  it  is  at  once  decom- 
posed, and  oxygen  is  evolved. 

Experiment  61. — Place  a  small  quantity  of  sodium  peroxide  in 
a  dry  flask  fitted  with  a  delivery  tube  and  a  stoppered  funnel,  as  in 
Fig.  48.  Allow  water  to  enter  the  flask  drop  by  drop  by  means  of 
the  funnel.  As  each  drop  falls  upon  the  powder,  a  brisk  action  is 

1  If  a  large  quantity  of  such  impurity  were  present,  such  as  would 
occur  if  the  black  oxide  of  manganese  were  adulterated  with  powdered 
coal,  it  would  give  rise  to  an  explosion  when  heated. 


Oxygen  from  Air. 


57 


noticed,  and  oxygen  is  rapidly  given  off,  which  may  be  collected 
in  the  usual  way. 

After  the  experiment,  dip  a  piece  of  turmeric  paper  into  the 
liquid  in  the  flask  ;  notice  that  the  paper  is  stained.  Also  dip  the 
fingers  into  it,  and  note  its  caustic  nature. 
Recall  the  solution  obtained  oy  the  action 
of  sodium  upon  water  (see  Exp.  43).  Here 
we  have  the  same  substance,  caustic  soda, 
formed  as  a  second  product  of  the  chemi- 
cal change. 


FIG.  48. 


(4)  Oxygen    from   Water. — Al- 
though water  is  the  commonest  com- 
pound of  oxygen,  we  very  seldom  em- 
ploy it  for  obtaining  this  gas,  because 
water    is    not    so    easily    decomposed 

as  many  other  oxygen  compounds.  There  are  very  few 
elements  which  have  a  sufficiently  strong  affinity  for  the 
hydrogen  of  the  water,  for  them  to  tear  it  away  from  the 
oxygen,  and  let  the  latter  go  free.  Under  certain  conditions, 
however,  the  element  chlorine  (also  a  gas)  is  able  to  do  this. 
Thus,  if  a  mixture  of  chlorine  gas  and  steam  be  strongly  heated 
by  being  passed  through  a  red  hot  tube,  the  chlorine  seizes  the 
hydrogen  of  the  water,  unites  with  it  to  produce  the  compound 
hydrogen  chloride  (or  hydrochloric  acid),  and  the  oxygen  in 
the  water  is  set  free. 

(5)  Oxygen  from  the  Air.— Oxygen  is  now  obtained  on 
a  manufacturing  scale  from  the  enormous  store  of  it  which  is 
present  in  the  air.     This  is  done  in  two  operations;   in  the 
first  the  atmospheric  oxygen  is  made  to  combine  with  some 
substance,  and  in  the  second  this  substance  is  again  decom- 
posed.    The   substance  employed   is   barium  oxide  (baryta). 
This  is  heated  in  iron  pipes  through  which  air  is  pumped  under 
a  slightly  increased  pressure.     Under  these  circumstances  the 
barium  oxide  combines  with  oxygen  from  the  air,  and  gives 
barium  dioxide,  while  the  other  chief  constituent  of  the  air, 
namely,  the  nitrogen,  passes  away.     Presently  the  pumps  are 
reversed,  and  a  partial  vacuum  is  produced  in  the  heated  pipes  ; 
this  causes  the  barium  dioxide  to  decompose,  changing  back  into 


5  8  Oxygen. 

the  original  barium  oxide,  and  giving  up  the  oxygen  it  had 
absorbed  from  the  air.  Therefore,  by  pumping  air  in  and  out 
of  these  heated  pipes  containing  baryta,  it  is  possible  to  obtain 
large  quantities  of  oxygen  very  rapidly.  This  method  of 
obtaining  oxygen  is  known  as  "Erin's  process." 

The  Properties  of  Oxygen. — From  the  various  speci- 
mens of  the  gas  which  have  been  prepared,  we  see  that  it  is 
colourless,  and  that  it  is  so  little  soluble  in  water  that  we 
observe  no  loss  while  collecting  it  in  the  pneumatic  trough. 
If  we  take  one  of  the  jars  and  apply  the  nose  to  the  gas,  we 
shall  find  that  oxygen  has  no  smell  or  taste.  This  is  exactly 
what  we  might  expect,  when  we  remember  that  the  air  we 
breathe  contains  a  large  proportion  (one-fifth)  of  free  oxygen. 

Test  for  Oxygen. — Oxygen  is  generally  recognized  and 
distinguished  from  other  gases  by  thrusting  into  a  jar  of  it  a 
chip  or  splinter  of  wood  which  has  been  lighted,  and  has  only 
a  glowing  spark  upon  it.  The  splinter  is  instantly  rekindled 
and  bursts  into  flame.  There  is,  however,  one  other  gas 
(namely,  nitrous  oxide)  which  behaves  in  a  similar  manner 
towards  a  glowing  splint.  Therefore  this  test  does  not  dis- 
tinguish oxygen  from  nitrous  oxide.  How  these  two  gases  are 
identified  is  explained  on  p.  202. 

Experiment  62. — Test  the  samples  of  gas  obtained  by  heating 
mercuric  oxide,  and  from  sodium  peroxide,  by  plunging  into  them 
a  splinter  of  wood  which  has  a  glowing  spark  upon  the  end.  The 
splinter  will  almost  immediately  rekindle.  Blow  it  out  again,  and, 
while  the  end  is  still  glowing,  thrust  it  once  more  into  the  gas.  It 
again  bursts  into  flame  as  before.  This  can  be  repeated  so  long  as 
sufficient  oxygen  is  left  in  the  jar. 

All  substances  which  are  capable  of  burning  in  the  air,  will 
burn  more  rapidly  and  with  increased  brilliancy  in  pure  oxygen. 
It  is  entirely  on  account  of  the  oxygen  present  in  the  atmo- 
sphere, that  substances  burn  in  the  air  at  all,  and  it  will  be 
evident  that  they  must  burn  more  quickly  in  oxygen  alone,  than 
in  oxygen  which  is  diluted  with  a  large  quantity  of  nitrogen,  as 
it  is  in  the  atmosphere.  We  may  prove  this  by  burning  a 
number  of  substances  in  the  gas  already  prepared. 


Properties  of  Oxygen.  59 

Experiment  63.— Charcoal  in  Oxygen.  Take  a  piece  of  char- 
coal about  the  size  of  a  hazel-nut,  and  fasten  it,  by  means  of  a  piece 
of  thin  copper-wire,  to  a  deflagrating  spoon.  Light  one  corner  of 
the  charcoal  in  a  gas-flame,  and  then  lower  the  spoon  into  a  jar  of 
oxygen.  Notice  that  the  charcoal  at  once  begins  to  burn  much 
more  brightly  than  it  did  before.  If  the  charcoal  used,  happens  to 
be  a  piece  made  from  the  bark  of  the  wood,  it  will  throw  off  a 
shower  of  brilliant  sparks  or  scintillations  as  it  burns  in  the  oxygen. 
When  the  charcoal  has  burnt  out,  remove  the  spoon,  and  pour  a 
little  water  into  the  jar.  Cover  it  with  a  glass  plate  and  keep  it  for 
a  further  experiment. 

Experiments^. — Sulphur  in  Oxygen.  Unscrew  the  small  cup 
from  a  deflagrating  spoon,  and  tie  a  small  bundle  of  asbestos  to  the 
end  of  the  wire  by  means  of  thin  copper  wire.  Then  melt  a  little 
sulphur  in  a  test-tube  and  dip  the  asbestos  into  it  so  as  to  get  it 
coated  over  with  the  sulphur.  Now  light  the  sulphur  upon  the 
asbestos  by  means  of  a  lamp  flame,  allow  it  to  burn  in  the  air  for  a 
moment  or  two  and  plunge  it  into  a  jar  of  oxygen.  Observe  the 
greatly  increased  brilliancy  of  the  flame  the  moment  it  comes  into 
the  oxygen.  Notice  that  the  burning  sulphur  produces  a  little 
smoke  or  fume  in  the  jar.  When  the  sulphur  has  burnt  out,  pour 
a  little  water  into  the  jar,  shake  it  up,  cover  the  jar  over  with  a 
glass  plate,  and  keep  it  for  a  subsequent  experiment. 

Experiment  65. — Phosphorus  in  Oxygen.  Take  a  piece  of 
phosphorus l  about  the  size  of  a  pea,  wipe  it  dry  with  blotting  paper, 
and  place  it  in  a  deflagrating  spoon.  Set 
fire  to  the  phosphorus  by  touching  it  with 
the  end  of  a  wire  which  has  been  just 
warmed  in  a  flame,  and  lower  the  spoon 
into  a  large  flask  filled  with  oxygen.  Ob- 
serve how  much  more  intense  is  the  light 
of  the  phosphorus  burning  in  oxygen  than 
burning  in  the  air.  Watch  the  experiment 
closely,  and  notice  that,  as  the  burning  goes 
on,  suddenly  the  light  becomes  almost  un- 
bearably dazzling,  and  then  quickly  dies 
down.  At  that  point  the  phosphorus  was 
rapidly  boiling. 

In  preparing  to  ao  tnis  experiment,  the 

wire  of  the  deflagrating  spoon  should  be  pushed  so  far  through  the 
metal  cap  that  the  cup  containing  the  phosphorus  reaches  nearly 

;  Remember  the  precautions  as  to  handling  phosphorus  given  on  p.  10. 


60  Oxygen. 

to  the  bottom  of  the  flask  as  shown  in  Fig.  49,  so  as  to  prevent  the 
flame  from  cracking  the  flask  at  its  shoulder. 

When  the  phosphorus  is  burnt  out,  remove  the  spoon,  add  a 
little  water,  and  shake  it  up  in  the  flask.  Notice  that  the  dense 
white  fumes  in  the  flask  rapidly  disappear ;  they  are  dissolved  by 
the  water. 

Experiment  66. — Pour  a  little  blue  litmus  solution  into  the  jars 
used  in  Exps.  63,  64,  65.  Notice  that  the  blue  solution  is  turned 
red  in  each  case ;  but  observe  that  in  the  jar  in  which  the  carbon 
was  burnt  the  red  colour  is  less  of  a  scarlet,  and  more  inclined 
to  purple  than  in  the  other  two  cases. 

Substances  which  have  the  power  of  turning  litmus  red, 
are  said  to  be  acid ;  therefore,  when  we  burn  carbon,  sulphur, 
and  phosphorus  in  oxygen,  and  add  water  to  the  products  of 
the  burning,  we  obtain  substances  which  are  acids. 

Experiment  67.— Sodium  in  Oxygen.  Put  a  small  fragment 
of  sodium  in  a  clean  dry  deflagrating  spoon,  and  heat  it  in  a  gas 
flame  until  the  sodium  begins  to  burn  in  the  air,  then  plunge  it 
into  a  cylinder  of  oxygen.  The  sodium  burns  very  brightly,  and 
produces  a  white  smoke.  The  product  of  the  combustion  is,  how- 
ever, in  this  instance  a  solid,  and  most  of  it,  therefore,  remains 
behind  in  the  spoon.  When  the  sodium  has  all  burnt,  remove  the 
spoon,  and  rinse  out  the  jar  with  a  little  water,  and  pour  the  liquid 
cut  into  two  small  beakers.  To  one  add  a  little  litmus,  and  notice 
that  it  is  not  turned  red.  To  the  other  add  a  little  litmus  which 
has  been  just  turned  red  by  a  single  drop  of  dilute  acid,  and  observe 
that  this  reddened  litmus  is  turned  back  to  its  original  blue  colour.1 

Substances  which  restore  the  blue  colour  to  reddened  litmus 
are  said  to  be  alkaline;  therefore,  when  sodium  is  burnt  in 
oxygen  and  the  product  dissolved  in  water,  we  obtain  an  alkali. 

Many  substances  which  will  not  burn  in  the  air,  will  burn 
readily  in  pure  oxygen. 

Experiment  68. — Iron  in  Oxygen.  In  order  to  burn  iron  in 
oxygen,  it  must,  like  the  phosphorus  or  the  sulphur  or  the  charcoal, 
be  first  lighted.  To  do  the  experiment  it  is  best  to  use  a  jet  of 

1  If  the  liquid  obtained  by  rinsing  out  the  jar  does  not  contain  enough 
of  the  product  from  the  burning  sodium  to  show  this  result,  the  main 
portion  which  was  left  in  the  spoon  may  be  used,  by  dissolving  it  off  in  a 
little  water. 


Burning  Iron  in  Oxygen. 


61 


oxygen  direct  from  a  store  of  the  gas  contained  either  in  a  gasholder 
or  better  in  a  metal  cylinder,  into  which  the  oxygen  has  been  pumped 
under  great  pressure. 

Connect  to  a  gas  reservoir  by  means  of  a  rubber  tube,  a  glass 
tube,  drawn  to  a  jet  at  one  end,  and  allow  a  stream  of  oxygen 
to  blow  a  spirit  lamp  flame  against  the  ends  of  a  bundle  of  fine  steel 
wires,  in  the  manner  shown  in  Fig.  50.  Almost  immediately  the 
tips  of  the  wires  get  red  hot,  and  begin  to  burn.  Now  remove 
the  lamp  and  the  metal  will  continue  to  burn  in  the  jet  of  oxygen, 
throwing  off  a  shower  of  brilliant  sparks.  Notice  that  the  end  of 
the  bundle  of  wires  melts,  and  drops  of  molten  matter  continually 
fall.  These  should  be  allowed  to  drop  into  an  iron  basin  of  water 
placed  ready  to  catch  them. 

In  the  absence  of  a  reservoir  of  oxygen,  fill  a  common  wide- 
mouthed  bottle  (such  as  a  pickle  bottle  or  glass  marmalade  jar)  with 


FIG.  50. 

oxygen,  leaving  about  two  inches  of  water  in  the  bottle.  Stick  into 
a  loosely  fitting  cork  a  bundle  of  steel  wire  or  a  straightened  piece 
of  watchspring,  the  end  of  which  has  been  tipped  with  sulphur. 
Light  the  sulphur  in  a  flame  and  quickly  plunge  the  wire  into  the 
bottle  of  oxygen.  The  burning  sulphur  will  set  fire  to  the  iron, 
which  will  go  on  burning  in  the  oxygen.  The  melted  product  of 
the  combustion  will  be  partially  quenched  by  the  layer  of  water  ; 
but,  in  spite  of  this,  it  will  probably  crack  the  bottle. 

Experiment  69. — Collect  some  of  the  solid  product  of  the 
burning  iron,  and  place  it  upon  a  moistened  litmus  paper.  Notice 
that  the  paper  is  not  turned  red. 

Redden  another  piece  of  litmus  paper  by  dipping  it  into  water 
containing  a  drop  or  two  of  acid,  and  place  some  of  the  substance 
upon  this.  The  blue  colour  is  not  restored  ;  therefore  the  product 
obtained  by  burning  iron  in  oxygen  is  neither  acid  nor  alkaline. 

In  each   case  the   product  obtained  when  a  substance  is 


62 


Oxygen. 


burnt  in  oxygen,  is  a  compound  of  the  thing  burnt  with  the 
oxygen.  The  process  of  burning  in  these  instances  is  there- 
fore nothing  more  than  the  rapid  combination  of  the  various 
substances  with  oxygen.  If  we  were  to  burn  carbon,  sulphur, 
phosphorus,  or  sodium  in  the  air,  and  examine  the  products 
obtained,  we  should  find  that  in  all  cases  they  were  the  same 
as  when  these  things  were  burnt  in  oxygen.  This  shows  that 
the  ordinary  process  of  combustion  in  the  air  is  also  the  rapid 
combination  of  substances  with  oxygen,  the  difference  being 
that  the  combination  is  not  so  rapid  as  in  pure  oxygen.  We 
can  easily  prove  that  it  is  the  oxygen  present  in  the  air  which 
enables  substances  to  burn  in  air,  by  removing  the  oxygen 
from  a  quantity  of  air,  and  trying  the  experiment  of  putting 
a  lighted  candle  into  the  air  that  is  left. 

Experiment  70. — Place  a  fragment  of  phosphorus  (wiped  dry) 
in  a  little  dish,  and  float  the  dish,  upon  water  in  a  pneumatic 
trough.  Set  fire  to  the  phosphorus,  and 
cover  it  with  a  cylinder  a  little  wider 
than  the  small  dish  as  in  Fig.  51.  The 
phosphorus  burns  in  the  confined  air 
in  the  cylinder,  and  combines  with  the 
oxygen  present,  forming  the  same  white 
fume  as  when  burnt  in  pure  oxygen. 
When  the  phosphorus  has  burnt  itself 
out,  allow  the  apparatus  to  stand  for  a 
short  time,  so  that  the  jar  may  cool,  and 
the  fumes  may  become  dissolved  by  the 
water.  Notice  that  the  water  has  risen 
in  the  cylinder,  taking  the  place  of  the 
oxygen  that  has  been  withdrawn.  Now 
slip  a  glass  plate  beneath  the  mouth 

of  the  cylinder  and  remove  it  from  the  trough.  Shake  the  air 
and  water  together  so  as  to  completely  dissolve  the  remaining  fume. 
Now  introduce  a  lighted  taper  or  candle  into  the  gas  in  the  jar,  and 
notice  that  the  flame  is  at  once  extinguished.  A  lighted  spill  of 
paper,  or  a  coal-gas  flame,  will  behave  in  a  similar  manner.  We 
have  withdrawn  all  the  oxygen  from  the  air  in  the  cylinder,  and 
the  gas  that  remains  (namely,  nitrogen)  will  not  support  the  com- 
bustion of  ordinary  burning  bodies. 

Oxides. — The   products   obtained   when   substances   are 


FIG.  51. 


Oxidation.  63 

burnt  in  oxygen  (either  pure  oxygen,  or  the  oxygen  of  the 
air)  are  called  oxides.  All  the  elements,  except  fluorine,  have 
been  made  to  combine  with  oxygen  either  directly  or  in- 
directly. 

We  have  seen  from  Exps.  66,  67,  and  68,  that  different 
oxides  have  very  different  properties  ;  thus,  some  combine 
with  water  to  yield  acids,1  while  others,  under  the  same  cir- 
cumstances, give  compounds  which  are  alkaline.  The  former 
of  these  are  called  acid-forming  oxides,  while  the  latter  belong 
to  a  class  known  as  basic  oxides.  (As  a  general  rule  the  non- 
metals  combine  with  oxygen  to  give  acid-forming  oxides,  while 
the  metals  yield  basic  oxides.  There  are,  however,  some 
oxides  of  both  non-metals  and  of  metals  which  are  neither 
acid-forming  nor  basic.) 

Hydroxides. —This  is  the  name  applied  to  the  com- 
pounds which  are  produced  when  oxides  combine  with  water. 
The  acid-forming  oxides  yield  acid  hydroxides,  while  those 
derived  from  basic  oxides  are  called  basic  hydroxides.  The 
term  acid  hydroxide,  however,  is  not  very  often  employed, 
as  these  compounds  are  included  in  the  class  of  substances 
called  acids ;  and  the  name  hydroxide  alone  is  more  usually 
used  to  denote  the  basic  hydroxides. 

Since  the  basic  hydroxides  are  derived  from  oxides  of 
metals,  they  consist  of  a  metal  combined  with  oxygen  and 
hydrogen  :  they  are  the  hydroxides  of  metals.  Thus,  sodium 
combined  with  oxygen  and  hydrogen  (or  sodium  oxide  com- 
bined with  water)  is  sodium  hydroxide ;  calcium  united  with 
oxygen  and  hydrogen  (or  calcium  oxide  united  with  water)  is 
calcium  hydroxide  ;  and  so  on. 

Oxidation. — The  process  of  combining  with  oxygen  is 
called  oxidation,  whether  the  action  takes  place  rapidly,  as 
when  substances  burn  in  oxygen,  or  whether  it  goes  on  slowly 
without  any  visible  signs  of  heat.  We  saw  in  Ex  p.  61,  the 
rapid  oxidation  of  the  metal  sodium,  but  the  same  process  of 
oxidation  goes  on  if  sodium  is  merely  exposed  to  oxygen,  even 
the  oxygen  of  the  air,  without  being  made  to  burn. 

1  For  further  development  of  this  subject,  see  p.  66. 


64  Oxygen. 

Experiment  71. — Take  a  fair  sized  piece  of  sodium  and  quickly 
cut  a  slice  off  it.  Notice  that  for  an  instant  the  freshly  cut  surface 
of  the  metal  looks  bright  and  silvery,  but  that  almost  immediately 
it  becomes  dull  and  tarnished.  It  quickly  gets  coated  with  a  white 
film  or  crust  which  consists  of  the  oxide  of  sodium.  The  tarnishing 
is  the  oxidation  of  the  metal,  and  if  the  piece  of  sodium  be  left 
exposed  to  the  air  it  is  soon  oxidized  right  through,  the  whole  piece 
being  changed  into  the  oxide. 

Other  metals  behave  in  a  similar  way  ;  thus,  when  bright 
iron  is  exposed  to  the  air,  we  know  that  it  soon  loses  its 
brilliant  surface,  and  becomes  coated  with  a  reddish  film  of 
rust.  This  rust  is  simply  the  oxide  of  iron,  formed  by  the 
slow  combination  of  the  metal  with  the  oxygen  of  the  air. 

Increase  of  Weight  by  Burning. — If  the  ordinary 
process  of  burning  is  simply  the  rapid  combination  of  the 
burning  substance  with  the  oxygen  of  the  air,  we  ought  to 
find  that  the  products  of  burning  actually  weigh  more  than 
the  material  that  is  burnt.  This  we  can  easily  put  to  the  test 
of  experiment. 

Experiment  72. — Place  a  heap  of  "  reduced  iron  " l  on  a  little 
iron  dish  or  tray,  and  balance  it  upon  a  small  pair  of  scales.  Then, 
by  means  of  a  lighted  taper,  set  the  heap  alight.  Observe  that  the 
black  powder  gradually  smoulders  right  through,  and  turns  to  the 
familiar  brownish  colour  of  iron  rust ;  but  note  also  that  as  it 
burns  the  mass  gains  in  weight,  for  the  scale  pan  on  which  it  is 
soon  begins  to  fall.  This  shows  that  the  rust  is  heavier  than  the 
iron  which  produced  it. 

In  this  experiment  the  only  product  of  the  burning  is  solid 
and  visible,  but  when  we  burn  a  candle  we  see  no  products. 
The  candle  simply  appears  to  waste  away,  leaving  nothing 
behind ;  by  burning  it  we  seem  to  have  completely  annihilated 
it.  Is  this  really  the  case  ?  or  can  it  be  that  the  products  of 
the  burning  of  a  candle  are  gases,  and  escape  unobserved  into 
the  air  ?  Let  us  try  the  following  experiment. 

1  Reduced  iron  is  readily  prepared  by  heating  oxide  of  iron  in  a  glass 
tube,  and  passing  a  stream  of  coal-gas  or  hydrogen  through  the  tube  until 
the  material  is  black. 


Respiration. 


Experiment  73 . — Take  a  piece  of  wide  glass  tube  (such  as  a 
lamp  chimney)  and  fill  the  upper  part  with  lumps  of  solid  caustic 
soda,  which  are  kept  in  position  by  first  hanging  a  false  bottom 
of  wire  gauze  into  the  tube  by  means  of  wires.  Fit  a  cork  into 
the  bottom,  on  which  is  fast- 
ened a  short  candle,  the  cork 
being  bored  with  holes  to 
allow  air  to  enter  (Fig.  52). 
Balance  this  arrangement 
upon  a  pair  of  scales,  and  then 
light  the  candle  and  quickly 
replace  .the  cork.  Again  note 
that  as  the  candle  burns  the 
apparatus  becomes  heavier. 
The  candle  in  the  act  of  burn- 
ing is  combining  with  oxygen, 
and  two  invisible  gaseous  pro- 
ducts are  formed,  namely, 
steam  and  carbon  dioxide, 
both  of  which  are  caught  or 
absorbed  by  the  lumps  of 
caustic  soda. 


FIG.  52. 


Matter  is  Indestruc- 
tible.— We  can  destroy  a 
candle  by  burning  it,  but  not  the  carbon  and  hydrogen  of  which  the 
candle  is  composed ;  these  live  on,  as  it  were,  but  in  different 
states  of  combination.  Similarly  if  we  burn  a  piece  of  sulphur 
it  disappears  from  view,  but  the  sulphur  is  not  destroyed,  only 
passed  into  combination  with  oxygen,  forming  a  compound 
which  is  gaseous  and  invisible. 

Respiration. — Not  only  is  oxygen  necessary  to  support 
the  combustion  of  ordinary  burning  bodies,  but  it  is  also 
indispensable  to  the  process  of  respiration.  If  an  animal  is 
placed  in  air  from  which  the  oxygen  has  been  removed,  or  in 
any  mixture  of  gases  which  does  not  contain  free  oxygen,  it 
quickly  dies.  All  animals  require  oxygen  to  breathe.  Respira- 
tion is,  in  fact,  a  process  of  oxidation ;  air  is  drawn  into  the 
lungs,  and  a  portion  of  the  oxygen  is  absorbed  by  the  blood. 
This  oxygen-laden  blood  (which  has  a  bright  red  colour) 
passes  throughout  the  body,  and  exerts  its  oxidizing  power 

F 


66  Oxygen. 

upon  certain  compounds  containing  carbon  which  have  to  be 
removed  from  the  system,  with  the  result  that  carbon  dioxide 
is  produced  (the  same  compound  as  is  formed  when  carbon 
is  burnt  in  oxygen,  or  when  a  candle  burns  in  the  air)  and  is 
exhaled  in  the  breath.  When  the  red  blood  has  thus  parted 
with  its  oxygen  it  becomes  of  a  dark  colour,  and  is  known 
as  venous  blood;  this,  travelling  back  to  the  lungs,  is  there 
once  more  charged  with  oxygen,  again  to  carry  it  to  all  parts 
of  the  body.  We  can,  by  a  simple  experiment,  show  that  the 
air  exhaled  from  the  lungs  contains  the  same  gas  as  is  formed 
when  a  gas  flame  or  a  candle  burns  in  the  air. 

Experiment  74. — Place  some  clear  lime-water  in  a  test-tube, 
and  by  means  of  a  glass  tube  dipping  into  it,  bubble  the  breath 
from  the  lungs  through  it.  Note  that  the  first  portions  of  breath 
produce  little  or  no  effect,  but  that  presently,  as  air  from  the  deeper 
parts  of  the  lungs  is  expelled,  the  clear  solution  quickly  becomes 
milky. 

And  we  can  also  show  that  the  air  so  expelled  has  been 
deprived  of  some  of  its  oxygen  by  the  following  experiment. 

Experiment  75. — Collect  in  a  cylinder  over  water  the  breath 
exhaled  by  one  single  long  expiration,  emptying  the  lungs  as  far 
as  possible  (see  Exp.  31,  Fig.  12).  Now  remove  the  cylinder  and 
lower  into  it  a  lighted  candle  on  a  wire.  Notice  that  the  flame  will 
be  extinguished.  [This  result  is  due  partly  to  the  presence  of  the 
carbon  dioxide,  as  well  as  to  the  smaller  amount  of  oxygen.] 

Acids. — There  is  a  class  of  compounds  which  chemists 
call  acids.  All  acids  have  a  sour  taste,  and  they  also  have  the 
power  of  reddening  a  solution  of  blue  litmus  (refer  back  to 
Exp.  66).  It  does  not  follow,  however,  that  all  sour  substances, 
or  all  that  will  redden  litmus,  are  acids.  This  is  by  no  means 
the  case,  for  we  know  of  many  things  which  have  a  sour  taste 
and  which  will  turn  litmus  from  blue  to  red,  but  which  do  not 
belong  to  the  class  of  substances  recognised  as  acids  ;  they  are 
acid,  in  the  sense  of  being  sour,  but  are  not  acids. 

The  question,  What  is  an  acid?  is  one  which  cannot  be 
answered  in  a  word,  and  chemists  are  not  agreed  as  to  what 
is  the  best  definition  to  be  given.  At  one  time  it  was  supposed 


Acids  and  Alkalies.  67 

that  all  acids  contained  oxygen,  that  this  element  was  the 
acidifying  principle  in  these  substances.  Indeed  the  very  name 
oxygen  means  the  acid  producer,  and  was  originally  applied  to 
this  element  by  Lavoisier  because  of  this  idea.  We  now  know 
that  this  belief  was  wrong,  for  we  have  many  acids  which  do 
not  contain  any  oxygen  in  their  composition.  At  the  present 
time  chemists  regard  the  element  hydrogen  as  a  necessary 
constituent  of  an  acid.  According  to  modern  notions,  a  com- 
pound which  does  not  contain  any  hydrogen  cannot  be  an 
acid.  This  of  course  does  not  mean  that  all  compounds 
containing  hydrogen  are  acids.  We  know,  for  example,  that 
water  is  a  compound  of  hydrogen,  but  water  is  not  an  acid,  it 
is  not  sour,  and  does  not  redden  litmus.  We  have  already 
learnt  that  sulphuric  acid  and  hydrochloric  acid  contain 
hydrogen,  because  we  obtained  this  element  from  both  of 
them  by  acting  upon  them  with  various  metals.  Can  we, 
therefore,  define  an  acid  as  a  compound  from  which  hydrogen 
can  be  expelled  by  the  action  of  certain  metals  ?  No,  because 
we  have  also  learnt  that  hydrogen  can  be  expelled  from  water 
by  means  of  certain  metals,  and  water  is  not  an  acid.  Perhaps 
the  best  definition  that  can  be  given  is  the  following,  an  acid 
is  a  compound  from  which  hydrogen  can  be  displaced  by  the 
hydroxide  of  a  metal.  The  word  displaced  here  only  means 
that  the  hydrogen  is  turned  out  of  the  compound,  and  not  that 
it  is  set  free,  or  liberated  as  gas. 

Alkalies. — This  name  is  applied  to  the  hydroxides  of  the 
alkali  metals  (of  which  sodium  and  potassium  are  the  most 
important)  and  to  ammonia.  When  sodium  is  burnt  in  air  or 
in  oxygen,  and  the  oxide  so  obtained  is  dissolved  in  water, 
sodium  hydroxide  is  produced  (Exp.  67).  We  have  learnt 
that  the  alkali  produced  in  this  experiment  has  the  property 
of  changing  the  colour  of  turmeric  from  yellow  to  brown,  and 
also  of  restoring  the  blue  colour  to  litmus  which  has  been 
reddened  by  an  acid. 

Experiment  76.— Make  a  dilute  solution  of  sodium  hydroxide 
(caustic  soda)  by  dissolving  a  small  piece  of  the  solid  in  water  in 
a  beaker.  Dip  a  turmeric  paper  into  it,  and  observe  the  stain. 
Make  a  dilute  sample  of  hydrochloric  acid  by  adding  a  little  strong 


68  Oxygen. 

acid  to  some  water  in  a  beaker.  Dip  a  turmeric  paper  into  this  ; 
note  that  no  stain  is  produced.  Add  to  the  acid  a  few  drops  of  a 
solution  of  litmus,  which  will  be  at  once  changed  from  blue  to  red. 
Now  add  the  dilute  alkali  very  gradually  to  the  acid,  and  watch 
the  effect  upon  the  reddened  litmus.  Where  the  alkali  meets  the 
acid  solution,  the  litmus  shows  a  blue  colour,  which,  however,  dis- 
appears again  on  gently  shaking  or  stirring.  Presently,  however, 
as  more  alkali  is  added,  the  red  liquid  turns  entirely  blue.  The 
solution  is  now  no  longer  acid,  but  alkaline.  Now  add,  drop  by 
drop,  a  little  more  of  the  dilute  acid,  stirring  the  solution  ;  stop  the 
moment  the  liquid  turns  red.  If  this  is  carefully  done,  one  drop 
of  the  alkali  will  now  restore  the  blue  colour.  Now  test  this 
solution  with  a  turmeric  paper,  and  note  that  no  brown  stain  is 
produced  ;  therefore  the  liquid  is  neither  add  nor  alkaline. 

Neutralization. — Exp.  76  teaches  us  that  when  we  mix 
an  acid  and  an  alkali,  each  one  destroys  the  other,  so 
that  a  point  can  be  arrived  at  when  the  solution  has  neither 
the  property  of  an  acid  nor  an  alkali.  Under  these  circum- 
stances we  say  that  the  liquid  is  neutral,  that  the  acid  has 
neutralized  the  alkali,  or  the  alkali  has  neutralized  the  acid. 
Now  we  must  ask  ourselves  what  has  become  of  the  acid  and 
the  alkali  ?  Are  both  still  present,  and  the  properties  of  each 
simply  masked  by  those  of  the  other  ?  In  other  words,  have 
we  simply  a  mixture  of  these  things  so  adjusted  that  the 
properties  of  the  one  just  balance  those  of  the  other,  or  has 
any  chemical  change  taken  place  between  the  acid  and  the 
alkali  resulting  in  new  compounds  which  happen  to  have  no 
effect  upon  either  turmeric  or  litmus  ?  Let  us  test  this  matter 
by  experiment. 

Experiment  77. — Take  some  of  the  dilute  acid,  the  alkali, 
and  the  neutral  mixture,  and  evaporate  each  separately  to  dryness 
in  small  dishes.  Note  first  that  the  acid  leaves  no  residue  behind  : 
that  is,  it  is  a  volatile  acid ;  therefore,  if  the  neutral  solution  is 
simply  a  mixture  of  the  two  substances,  when  it  is  evaporated 
down,  the  acid  ought  to  volatilize  along  with  the  steam  and  leave 
the  alkali,  in  which  case  the  residue  in  the  dish  containing  the 
alkali,  and  that  in  the  one  containing  the  neutral  solution  would 
be  the  same.  Are  they  the  same?  Observe  that  they  appear 
different.  Touch  each  with  a  moist  turmeric  paper,  the  one  is 


Salts.  69 

strongly  alkaline,  while  the  other  is  still  perfectly  neutral,  therefore 
they  are  different.  The  residue  being  neutral  shows  that  by 
evaporating  the  solution,  the  acid,  which  had  been  added,  has  not 
been  driven  off ;  it  cannot,  therefore,  have  been  simply  mixed  with 
the  alkali,  but  must  have  entered  into  chemical  union  with  it. 
Now,  with  the  little  finger  bring  a  little  of  the  neutral  residue  on  to 
the  tip  of  the  tongue  ;  the  familiar  taste  of  the  substance  will  also 
prove  that  we  have  here  a  compound  which  is  quite  different  from 
either  sodium  hydroxide  or  hydrochloric  acid. 

We  learn,  therefore,  that  when  acids  and  alkalies  neutralize 
one  another  they  enter  into  chemical  combination  with  each 
other,  forming  new  compounds. 

Bases. — A  large  number  of  other  substances  besides  the 
alkalies,  are  capable  of  neutralizing  acids.  Many  of  these 
substances  resemble  the  alkalies  in  having  the  power  of 
restoring  the  blue  colour  to  reddened  litmus,  and  in  imparting 
a  brown  colour  to  turmeric;  they  are,  therefore,  said  to  be 
alkaline  in  character,  or  to  possess  an  alkaline  reaction.  Some 
chemists,  indeed,  extend  the  name  alkali  so  as  to  include  many 
of  these  compounds.  The  term  base  has  long  been  in  use  to 
denote  any  substance  which  is  capable  of  neutralizing  an  acid, 
and  it  therefore  includes  the  alkalies.  For  the  most  part 
bases  are  compounds  of  metals;  being  either  the  oxides  or 
hydroxides  of  metals.  Ammonia,  however,  which  contains  no 
metal,  being  only  a  compound  of  nitrogen  and  hydrogen,  is 
also  included  in  this  class,  and  so  are  many  organic  com- 
pounds which  are  not  compounds  of  metals.  For  our  present 
purpose,  however,  we  may  define  bases  as  certain  oxides  and 
hydroxides  of  metals,  which  are  able  to  neutralize  acids,  and 
include  ammonia. 

Salts. — The  compounds  that  are  produced  when  acids 
and  bases  combine  together  are  called  salts.  We  have  seen 
by  Exp.  77  that  when  hydrochloric  acid  and  sodium 
hydroxide  neutralize  each  other  the  substance  we  familiarly 
call  "salt"  is  produced.  "Salt"  is  one  of  the  commonest 
and  most  important  of  all  salts,  as  well  as  being  one  that  has 
been  longest  known  to  man ;  and  the  compounds  belonging  to 
this  class  are  called  salts,  because  of  a  general  similarity  many 


70  Oxygen. 

of  them  bear  to  "salt,"  or  common  salf,  as  it  is  termed.  No 
satisfactory  definition  of  a  salt  can  be  given,  because  chemists 
are  not  all  agreed  as  to  exactly  what  shall  be  included  in  this 
class  of  compounds.  Some  regard  acids  themselves,  as  being 
salts  of  hydrogen.  At  this  stage  we  will  consider  salts  as 
being  the  compounds  that  are  produced  by  the  interaction  of 
acids  with  bases. 

EPITOME. 

Oxygen  was  discovered  in  1774  by  Priestley,  by  heating  mercuric1 
oxide.  It  is  the  most  abundant  of  all  the  elements.  Occurs  un- 
combined  in  the  air.  It  is  usually  obtained  by  heating  potassium 
chlorate :  more  readily  if  the  chlorate  be  mixed  with  manganese 
dioxide.  Oxygen  is  also  given  off  when  certain  peroxides,  such  as 
manganese  dioxide,  or  barium  dioxide,  are  strongly  heated  ;  the 
latter  substance  is  used  in  Erin's  process.  Oxygen  can  be  obtained 
from  water  by  heating  steam  and  chlorine  together. 

Oxygen  is  a  colourless,  tasteless,  and  inodorous  gas  ;  supports 
combustion  energetically,  and  is  necessary  to  life.  Oxygen  is 
slightly  soluble  in  water ;  100  volumes  of  water  can  dissolve  4 
volumes  of  oxygen.  Fish  depend  upon  this  dissolved  oxygen  for 
their  supply  of  this  gas  for  respiration  ;  they  cannot  breathe  free 
air,  neither  can  they  use  the  oxygen  which  is  a  chemical  constituent 
of  water. 

The  products  obtained  by  burning  things  in  oxygen,  or  in  air, 
are  oxides.  Some  of  these  are  acid-forming,  others  are  basic 
oxides.  Similar  compounds  are  formed  when  the  same  elements 
combine  slowly  with  oxygen,  the  process  then  being  called  oxidation. 
The  products  obtained  by  burning  or  by  oxidation  are  heavier  than 
the  original  substance  that  is  burnt ;  burning  is  only  matter  under- 
going change,  and  not  matter  being  destroyed. 

Acids  are  compounds  containing  hydrogen  which  can  be  dis- 
placed from  the  compound  by  the  hydroxide  of  a  metal.  They  are 
sour  or  acid  to  the  taste,  and  will  redden  blue  litmus. 

Bases  are  compounds  which  will  neutralize  acids.  They  are 
mostly  certain  oxides  and  hydroxides  of  metals.  Ammonia  is  also 
a  base.  Those  that  are  soluble  in  water  are  alkaline  ;  they  restore 
the  blue  colour  to  reddened  litmus,  and  give  a  brown  stain  to 
turmeric. 

Salts  are  the  compounds  produced  by  the  union  of  acids  and 
bases. 


Oxygen.  7 1 

Reactions  for  oxygen  x — 

(1)  From  mercuric  oxide  HgO  =  Hg  +  O 

(2)  „      potassium  chlorate  KC1O3  =  KC1  +  30 

(3)  „      sodium    peroxide  and  water    Na2O2  4-  H2O  =  2NaHO 

+  0 

(4)  „      water  by  the  action  of  chlorine  H2O  +  C12  =  2HC1  +  O 

(5)  ,-,      air,  by  Erin's  process  — (0)  BaO  +  O  =  BaO2 

(&)  BaO2  =  BaO  +  O 

Combustions  in  oxygen — 

Carbon  C  +  O2  =  CO2 
Sulphur  S  +  O2  =  SO2 
Phosphorus  aP  +  $O  =  P2O5 
Iron  3Fe  +  40  =  Fe3O4 

1  These  will  be  explained  later  on';  the  student  can  pass  them  over  at 
this  stage. 


CHAPTER   VIII. 

WATER. 

WE  have  learnt  by  Exp.  54  that  when  hydrogen  burns  in 
the  air,  water  is  formed ;  and  as  other  elements  on  burning 
in  the  air  yield  their  oxides,  so  we  might  expect  that  hydrogen 
would  do  the  same,  and  that  water  is  an  oxide  of  hydrogen. 
First  let  us  make  an  experiment  in  order  to  collect  the  liquid 
which  is  formed,  and  examine  it. 

Experiment  78. — Burn  a  small  flame  of  hydrogen  under  the 
open  end  of  a  bent  glass  tube,  arranged  as  in  Fig.  53.  A  liquid 
quickly  condenses  on  the  long  neck  and  runs  down  into  the  flask, 
and  in  about  half  an  hour  a  considerable  quantity  will  be  collected. 

One  property  by  which  we  can  recognize  water  from  other 
liquids  we  have  learnt  from  Exp.  44,  and  that  is  its  behaviour 
towards  the  metal  potassium. 

Experiment  79. — Pour  the  liquid  collected  in  Exp.  78  into  a 
test-tube,  and  drop  into  it  a  small  bit  of  potassium  ;  if  the  metal 
takes  fire,  as  it  did  in  Exp.  44,  we  conclude  that  the  liquid  is  water. 
[Other  tests  by  which  we  can  recognize  water  we  shall  learn  later  on.] 

As  the  air  does  not  consist  of  oxygen  alone,  we  cannot 
say  certainly  from  this  experiment  that  water  is  a  compound 
of  hydrogen  and  oxygen  only.  This  point  might  be  settled  by 
burning  the  hydrogen  in  pure  oxygen.  It  was  indeed  first 
settled  for  us  by  Cavendish,  in  1781. 

Cavendish  used  a  strong  glass  vessel,  similar  to  E,  Fig.  54, 
which  had  a  stopcock  at  the  bottom,  and  was  closed  at  the 
top  with  a  stopper  through  which  two  wires  were  fixed,  so  that 


Water. 


73 


he  could  send  an  electric  spark  into  the  gases  it  contained. 
He  pumped  all  the  air  out  of  this  vessel  with  an  air  pump,  and 
then  attached  it  to  the  bell  jar,  B,  containing  a  mixture  of 
hydrogen  and  oxygen.  On  opening  the  taps  the  gases  of 
course  entered  into  the  vacuous  tube.  The  stopcocks  were 
then  closed,  and  a  spark  from  an  electrical  machine  was  made 
to  pass  between  the  two  wires.  This  of  course  exploded  the 


r 


FIG.  53. 


FIG.  54. 


mixture  of  hydrogen  and  oxygen,  but  the  vessel  being  very 
strong  did  not  burst.  After  the  two  gases  had  thus  combined, 
there  was  a  minute  quantity  of  water  in  the  tube. 

By  again  filling  the  vessel  with  more  of  the  gas  and  ex- 
ploding the  mixture,  the  quantity  of  water  increased,  until  by 
repeating  the  experiment  several  times,  enough  of  the  liquid 
was  collected  to  prove  that  it  was  actually  pure  water.  As 
there  was  nothing  besides  hydrogen  and  oxygen  present  in  the 
gases,  this  proved  that  water  was  composed  simply  of  these 
two  elements. 

But  Cavendish  did  more  than  merely  prove  this,  he  also 


74 


Water. 


taught  us  the  proportions  in  which  these  two  gases  combined 
together  to  produce  water.  If,  for  instance,  the  mixture  of 
gases  in  the  bell  jar  contained  equal  measures  of  hydrogen  and 
oxygen,  he  found  that  after  the  explosion  there  was  always 
some  oxygen  left  over  in  the  vessel  E.  But  if  he  mixed  the 
gases  exactly  in  the  proportion  of  two  measures  of  hydrogen  to 
one  of  oxygen,  then  there  was  no  gas  left  in  E  after  the  ex- 
plosion ;  it  entirely  disappeared,  leaving  a  vacuum  in  the  vessel. 
Therefore  Cavendish  made  the  discovery  that  water  was  a 
compound  of  hydrogen  and  oxygen  only,  and  that  these  two 
elements  combined  in  the  proportion  of  tivo  volumes  of  hydrogen  to 
om  volume  of  oxygen. 

We  can  make  hydrogen  and  oxygen  combine  together  by 
what  are  called  indirect  methods,  and  thus  get  additional  proof 
that  water  is  composed  of  these  two  elements  only. 


FIG.  55. 


Experiment  80. — Roll  a  piece  of  fine  copper  gauze  into  short 
compact  cylinders,  c,  Fig.  55,  and  place  them  in  a  piece  of  hard 
glass  tube  (combustion  tube}.  To  one  end  of  this  tube  connect  a 
small  Wurtz  flask,  in  the  manner  shown  at  w,  Fig.  55,  and  to 
the  other  attach  a  small  apparatus  for  generating  oxygen  from  potas- 
sium chlorate.  Heat  the  combustion  tube  with  a  long  flat  flame 


Formation  by  Synthesis.  75 

until  the  copper  is  red  hot,  and  then  send  a  slow  stream  of  oxygen 
over  it  by  gently  heating  the  potassium  chlorate.  Notice  that  as 
the  oxygen  passes  over  the  copper,  the  latter  becomes  black.  It  is 
gradually  combining  with  the  oxygen  and  is  being  converted  into 
copper  oxide,  which  is  black.  After  some  little  time  disconnect  the 
oxygen  generator,  and  replace  it  by  an  apparatus  for  generating 
hydrogen  from  zinc  and  sulphuric  acid,  and  pass  a  slow  stream  of 
hydrogen  through  the  tube.1  Observe  that  almost  at  once  some 
moisture  condenses  in  the  Wurtz  flask,  and  as  the  experiment  goes 
on,  more  and  more  water  collects  in  the  little  receiver.  Where  is 
this  water  coming  from  ?  Notice  that  the  black  copper  oxide  is 
gradually  changing  back  again  into  bright  metallic  copper.  The 
hydrogen  has  taken  away  the  oxygen  which  the  copper  had 
combined  with,  and  has  united  with  it  to  form  water. 

In  this  experiment  therefore  we  have  caused  hydrogen  and 
oxygen  to  unite,  by  first  combining  the  oxygen  with  copper, 
and  then  allowing  hydrogen  to  deprive  the  copper  oxide  of  the 
oxygen,  whereby  the  copper  oxide  was  again  reduced  (that  is, 
the  copper  was  deprived  of  the  element  with  which  it  had 
combined,  and  was  restored  to  its  former  state  of  metallic 
copper),  and  water  was  formed.  This  experiment  is  a  very 
important  one,  and  we  shall  return  to  it  again  later.  In 
Cavendish's  experiment,  and  in  Exp.  78,  water  was  obtained 
by  the  direct  combination  of  its  constituent  elements; 
by  Exp.  80,  it  was  formed  by  the  indirect  union  of  oxygen 
and  hydrogen  ;  but  whether  directly  or  indirectly,  the 
compound  was  built  up,  so  to  speak,  from  the  elements  of 
which  it  is  composed.  This  method  of  proceeding  is  called 
synthesis. 

There  is  another  way  by  which  we  can  find  out  the  com- 
position of  a  compound  like  water,  which  is  exactly  the  opposite 
of  synthesis.  Instead  of  taking  the  constituents  of  the  com- 
pound and  making  them  combine,  we  can  take  the  compound 
and  decompose  it  into  its  constituents.  This  method  is  known 
as  analysis. 

1  A  convenient  piece  of  apparatus,  easy  to  make,  by  means  of  which  a 
regulated  stream  of  hydrogen  can  be  obtained  at  will,  is  described  on  p.  246, 
Fig.  106.  Put  granulated  zinc  in  one  tube  and  dilute  sulphuric  acid  in 
the  other. 


Water. 


For  instance,  in  Exps.  45,  48,  we  decomposed  water  by 
sodium,  and  by  iron,  and  got  out  of  the  water  one  of  its 
constituents,  namely,  the  hydrogen. 

By  the  action  of  chlorine  upon  water  (as  described  on 
p.  57),  the  water  was  again  decomposed,  and  the  other  con- 
stituent, namely,  the  oxygen,  was  obtained.  These  are 
processes  of  analysis. 

Water  can  also  be  decomposed  by  means  of  an  electric 
current  from  a  battery. 

Experiment  8 1.— Take  an  ordinary  wide  mouthed  bottle,  and  cut 
it  in  half.  (This  may  be  done  by  first  making  a  small  scratch  on  the 
glass  with  a  file,  and  then  touching  the  spot  with  a  red  hot  wire. 

The  glass  will  then  crack  at  the  file 
mark,  and  the  crack  can  be  made  to 
travel  right  round  the  bottle  by 
slowly  drawing  the  hot  wire  along 
just  in  front  of  the  crack.)  Fit  into 
the  neck  a  cork  through  which  two 
short  pieces  of  platinum  wire  have 
been  pushed.  The  apparatus  is  then 
supported  as  shown  in  Fig.  56,  and 
nearly  filled  with  water,  to  which  a 
little  sulphuric  acid  has  been  added.1 
Now  attach  the  two  wires  from  the 
battery  to  the  two  platinum  wires, 
and  notice  that  bubbles  of  gas  at 
once  make  their  appearance  on  each 
of  the  wires  in  the  acid  water.  Fill 
a  short  stout  test-tube  with  the  dilute 
acid  and  invert  it  over  the  two  wires 
in  the  basin,  and  collect  the  gas  that 
is  evolved.  When  the  tube  is  full, 
remove  it  and  apply  a  lighted  taper  to  the  gas.  If  it  is  hydrogen 
it  will  be  recognized  by  the  characteristic  flame ;  if  oxygen,  we 
shall  see  the  taper  flame  burn  more  brightly.  Note,  however,  that 
the  gas  does  not  answer  to  either  of  these  tests,  but  that  it  explodes 
with  a  sharp  crack.  This  shows  that  we  have  both  oxygen  and 
hydrogen  mixed  together. 

1  Pure  water  will  not  conduct  electricity,  hence  sulphuric  acid  is  added 
to  it. 


FIG.  56. 


The   Volume  Composition  of  Water.  77 

Experiment  82. — Now  collect  the  gas  from  each  wire  separately 
by  inverting  a  separate  tube  over  each  wire,  using  two  test-tubes  of 
the  same  size.  Notice  that  gas  collects  more  quickly  in  one  tube 
than  in  the  other.  As  soon  as  the  tube  which  fills  quickest  is  full 
down  to  the  water  level,  stop  the  experiment  by  disconnecting  one 
wire  from  the  battery.  Note  that  the  other  tube  is  only  half  full. 
Now  test  the  gas  in  each  tube,  and  find  that  the  gas  in  smallest 
quantity  will  rekindle  a  glowing  splint  of  wood ;  it  is  therefore 
oxygen  ;  while  the  other  burns  with  the  characteristic  flame  of 
hydrogen. 

These  two  experiments  prove  by  analysis  that  water 
consists  of  oxygen  and  hydrogen ;  and  also  that  it  contains 
these  two  gases  in  the  proportion  of  two  volumes  of  hydrogen 
to  one  volume  of  oxygen.  They  therefore  confirm  Cavendish's 
synthetical  experiments. 

When  oxygen  and  hydrogen  unite  (as  in  Cavendish's  ex- 
periment), the  volume  occupied  by  the  water  that  is  formed  is 
so  extremely  minute,  that  it  is  scarcely  measurable  unless  very 
large  volumes  of  the  two  gases  are  used.  We  cannot,  therefore, 
make  any  comparison  between  the  volume  of  the  gas  and 
that  of  the  liquid  water  that  is  produced.  But  if,  instead  of 
letting  the  water  condense  to  the  liquid  state,  we  were  to  make 
the  experiment  at  a  high  temperature,  so  that  the  water  was 
kept  in  the  condition  of  steam,  then  we  could  find  out  the 
relation  between  the  volume  of  the  mixed  gases,  oxygen  and 
hydrogen,  and  the  volume  of  the  gaseous  compound,  steam. 

This  experiment  is  made  in  the  following  way.  Two 
measures  of  hydrogen  and  one  of  oxygen  are  put  in  the  closed 
limb  of  the  U  tube  (Fig.  57),  which  is  divided  into  three 
equal  measures  by  rings  on  the  glass.  The  tube  is  then  heated 
by  boiling  some  amyl  alcohol  in  a  flask  and  sending  the  hot 
vapour  through  the  outer  glass  tube  which  surrounds  the  tube 
containing  the  gases.  The  vapour  leaves  this  "jacket  tube" 
by  the  small  pipe  at  the  bottom,  and  is  again  condensed  and 
collected.  When  the  whole  apparatus  is  hot,  the  mercury  is 
made  level  in  both  limbs,  and  the  gas  exactly  occupies  the 
three  measures.  Now  an  electric  spark  from  an  electrical 
machine  is  passed  between  two  platinum  wires  which  are 


78  Water. 

sealed  into  the  closed  end  of  the  tube  containing  the  gases 
This  causes  the  oxygen  and  hydrogen  to  combine  (which  the} 
do  with  some  violence,  so  that  precautions  are  taken  to  prevent 
the  mercury  from  being  blown  out  at  the  open  end  of  the  U 
tube)  and  form  water,  which,  however,  cannot  condense  to  the 
liquid  state,  but  remains  as  steam,  on  account  of  the  high 
temperature  of  the  vapour  surrounding  the  tube.  The  first 
thing  noticed  after  the  explosion  of  the  gases,  is  that  the 
volume  is  altered,  for  the  mercury  has  risen  in  the  tube. 


FIG.  57. 

Once  more,  therefore,  the  mercury  must  be  made  level  in  both 
limbs,  by  pouring  more  into  the  open  side ;  and  when  this  is 
done,  it  is  seen  that  the  gas  in  the  tube  now  occupies  exactly 
two  of  the  measures.  That  is  to  say,  two  volumes  of  hydrogen 
and  one  of  oxygen,  three  volumes  of  the  mixed  gases,  form  two 
volumes  of  steam  or  gaseous  water. 

The  Properties  of  Water. — Water  is  familiar  to  us  as 
a  colourless  and  tasteless  liquid.  When  we  look  through  a 
considerable  depth  of  it,  however,  it  is  seen  to  be  possessed 
of  a  bluish-green  colour.  We  know  that  when  water  is  cooled 


Point  of  Maximum  Density  of   Water.  79 

to  a  certain  temperature  it  solidifies  or  freezes ;  and  when 
heated  to  a  particular  temperature  it  boils.  One  method, 
therefore,  by  which  we  are  able  to  distinguish  water  from  any 
other  colourless  liquid  is  to  ascertain  its  freezing  and  boiling 
points.  The  freezing  point  of  water  is  o°  on  the  centigrade 
scale,  while  its  boiling  point  is  100°  (see  p.  88). 

As  water  is  gradually  cooled  down,  like  other  substances 
it  shrinks  in  volume.  This  shrinking  or  contraction  goes 
on  steadily  until  the  temperature  is  just  within  4°  of  the 
freezing  point.  After  this  temperature  is  passed,  the  water, 
instead  of  continuing  to  contract,  actually  expands  again,  just 
as  though  it  were  being  warmed.  Therefore,  if  we  take  some 
water  having  a  temperature  of  4°  C.,  it  will  expand  whether 
we  heat  it  or  cool  it.  At  this  particular  temperature,  water  is 
more  dense  than  at  any  other  point ;  4°  C.  is,  therefore,  called 
its  point  of  maximum  density. 

The  expansion  that  water  suffers  on  being  cooled  from 
4°  to  o°  (without  freezing)  is  extremely  slight;  1000  cc.  of 
water  measured  at  4°,  when  cooled  to  o°,  will  become  ex- 
panded to  1000-13  cc-  But  tnis  expansion,  although  appa- 
rently so  trifling,  plays  an  important  part  in  nature.  When 
lakes  and  ponds  are  exposed  to  the  cold  winter  winds,  the 
surface  water  becomes  cooled,  and  consequently  contracts  and 
becomes  denser.  It  therefore  sinks  to  the  bottom,  and  fresh 
layers  come  to  the  surface  to  be  cooled  in  their  turn,  and  also 
to  sink.  In  this  way  a  circulating  movement  is  set  up,  which 
continues  until  the  temperature  of  the  whole  mass  of  water 
has  fallen  to  4°  C.  When  this  state  is  reached,  any  further 
cooling  of  the  surface  expands  the  water  on  the  top.  It  there- 
fore becomes  less  dense,  and  so  remains  as  a  colder  layer 
floating  upon  the  surface,  until  at  last  it  solidifies  to  a  thin 
film  of  ice,  which  then  protects  the  water  beneath  from  contact 
with  the  cold  wind.  If  water  continued  steadily  to  contract 
with  cold  until  the  freezing  point,  this  circulating  movement 
would  not  stop  at  4°  C.,  but  would  go  on  until  the  whole  body 
of  water  in  lakes  and  ponds  had  reached  o°,  when  ice  would 
begin  to  form  at  the  bottom  as  well  as  the  top,  and  the  entire 
mass  of  water  would  verv  soon  become  solidified  throughout. 


8o  Water. 

When  water  actually  freezes,  it  suddenly  expands  very  con- 
siderably, 10  volumes  of  water  becoming  n  (nearly)  volumes 
of  ice.  Ice  is  therefore  lighter  than  water.  10  cc.  of  water 
weigh  10  grams,  but  n  cc.  of  ice  only  weigh  10  grams.  Con- 
sequently ice  floats  on  water. 

The  force  which  is  exerted  by  the  expansion  when  water 
freezes  is  very  great.  Thus,  if  a  strong  iron  bottle  be  filled 
entirely  with  water,  and  the  mouth  securely  closed,  and  then 
the  bottle  be  exposed  to  a  freezing  temperature,  the  expansion 
will  cause  the  bottle  to  burst.  It  is  owing  to  this  that  water- 
pipes  burst  in  frosty  weather.  The  pipe  splits,  or  bursts, 
at  the  moment  the  water  freezes,  but  it  is  only  when  the  ice 
melts  again  that  the  damage  to  the  pipe  is  revealed  ;  hence 
the  mistake  is  often  made  of  supposing  that  the  thaw  caused 
the  pipe  to  burst. 

We  say  that  water  boils  at  a  temperature  of  100°  C. ;  but, 
to  be  exact,  we  must  add  under  a  pressure  0/760  mm.  (see 
p.  92).  The  actual  temperature  at  which  water,  or  any  ether 
liquid,  boils,  depends  entirely  upon  the  pressure.  This  is  very 
easy  to  prove.  If  we  take  a  quantity  of  water  in  a  beaker, 
scarcely  hotter  than  can  be  borne  by  the  hand,  therefore  far 
short  of  boiling,  and  place  it  under  the  receiver  of  an  air- 
pump,  and  reduce  the  pressure  by  quickly  pumping  out  some 
of  the  air,  we  shall  see  the  water  begins  to  boil  violently, 
although  obviously  it  cannot  be  getting  any  hotter.  In  a 
vacuum  water  can  be  made  to  boil  even  at  the  freezing  tem- 
perature, therefore  boiling  has  not  necessarily  anything  to  do 
with  being  hot.  In  ordinary  life  we  are  in  the  habit  of  associating 
the  two  ideas  together,  so  that  the  very  word  boiling  is  almost 
synonymous  with  being  hot;  hence  the  common  expression 
"boiling  hot."  In  reality  the  expression  "boiling  cold"  is 
quite  as  correct,  for  many  boiling  liquids  are  colder  than  ice, 
and  colder  even  than  the  coldest  arctic  regions. 

The  Solvent  Power  of  Water. — Water  is  able  to  dis- 
solve a  great  many  substances,  and  chemists  make  use  of  this 
property  in  many  different  ways.  Thus  substances,  which  in 
the  solid  state  are  incapable  of  acting  chemically  upon  each 
other,  will  often  combine  if  either  one  or  both  of  them  are 


The  Hardness  of    Water.  8 1 

first  dissolved  in  water,  hence  a  large  number  of  chemical 
operations  (especially  analytical)  are  conducted  with  solutions 
in  water.  The  solvent  power  of  water  is  also  made  use  of 
for  separating  substances  which  are  more  soluble,  from  others 
which  are  either  slightly  soluble  or  are  altogether  insoluble 
(see  Exps.  20  and  27). 

Many  gases  are  soluble  in  water;  some  to  a  very  great 
extent,  others  only  very  slightly.  Thus,  i  cc.  of  water  at  o° 
will  dissolve  1148  cc.  of  ammonia  gas,  but  only  0*048  cc.  of 
oxygen.  In  all  cases,  gases  are  less  soluble  in  hot  than  in  cold 
water ;  thus  at  20°  i  cc.  of  water  can  only  dissolve  680  cc. 
of  ammonia,  and  0*028  cc.  of  oxygen.  Therefore,  if  a  solution 
of  a  gas  in  water  be  boiled,  all  the  gas  is  expelled. 

Natural  Waters. — Owing  to  its  great  solvent  powers, 
perfectly  pure  water  is  never  found  in  nature,  and,  indeed,  can 
only  be  obtained  in  the  laboratory  by  taking  great  precautions. 
The  purest  form  of  natural  water  is  rain  water,  but  even  this 
dissolves  the  gases  from  the  air,  and  is  also  contaminated  with 
impurities  that  are  present  in  the  air.  The  moment  rain 
touches  the  ground,  it  begins  to  dissolve  the  soil  or  rocks  upon 
which  it  falls  and  through  which  it  percolates,  and  it  gradually 
becomes  more  and  more  impure  as  it  finds  its  way  to  river 
and  ocean. 

Absolutely  pure  water  when  evaporated  to  dryness  in  a 
platinum  vessel  leaves  no  residue.  If  boiled  down  in  a  glass 
vessel  it  slightly  dissolves  the  glass.  Sea-water  contains  more 
dissolved  impurities  in  the  form  of  various  salts  than  other 
natural  waters.  Thus,  in  1000  grams  (i  litre)  of  the  water 
of  the  British  Channel,  there  are  35*25  grams  of  dissolved 
salts,  which  remain  as  a  residue  when  this  volume  of  the  water 
is  evaporated  to  dryness.  The  greater  part  of  this  residue, 
viz.  27  grams,  is  common  salt.  The  solid  dissolved  sub- 
stances in  good  drinking  water  average  from  0*437  to  0*03 
grams  in  i  litre. 

Hardness  of  Water. — Certain  of  the  dissolved  solids  in 
natural  waters  impart  the  property  known  as  hardness.  These 
salts  are  chiefly  the  carbonate  and  sulphate  of  lime.  When 
hard  water  is  boiled,  if  its  hardness  is  due  to  the  presence  of 

G 


82  Water. 

dissolved  carbonate  of  lime,  it  becomes  soft,  because  the 
carbonate  of  lime  is  precipitated  as  a  solid,  which  in  time 
collects  on  the  sides  of  the  vessel  (boiler  or  kettle),  and  pro- 
duces the  furring  which  is  so  common.  Sulphate  of  lime  is 
not  thrown  out  of  solution  by  boiling ;  therefore,  if  the  hardness 
is  due  to  this  salt,  boiling  the  water  will  not  soften  it.  On 
this  account,  hardness  caused  by  carbonate  of  lime  is  called 
temporary  hardness  ;  and  that  due  to  sulphate  of  lime  is 
termed  permanent  hardness.  If  a  sample  of  water  contains 
both  these  compounds,  then  it  loses  only  a  part  of  its  hard- 
ness, viz.  the  temporary  hardness,  on  boiling,  while  the 
permanent  hardness  remains. 

Water  of  Crystallization. — When  salts  are  crystallized 
from  solution  in  water  (as  in  Exp.  26),  it  often  happens  that 
some  of  the  water  solidifies  along  with  the  salt.  It  is  as 
though  the  particles  of  the  salt  were  unable  to  build  them- 
selves up  into  the  regular  shapes,  we  call  crystals,  without  the 
aid  of  some  particles  of  water  to  hold  them  together ;  just  as 
bricks  cannot  be  built  up  into  an  edifice  without  the  aid  of 
mortar.  If  the  mortar  were  to  be  removed  from  a  building, 
the  bricks  would  all  tumble  down  in  a  heap ;  and  if  we  remove 
the  solidified  water  from  such  a  crystal,  it  will  fall  down  to 
a  powder.  Water  which  is  thus  held  by  a  crystal,  and  which 
is  necessary  to  its  existence  as  a  crystal,  is  called  water  of 
crystallization. 

Experiment  83. — Take  a  few  crystals  of  copper  sulphate  (blue 
vitriol}  and  gently  heat  them  in  a  dish,  either  over  a  small  flame 
or  better  in  an  oven.  Notice  that  gradually  the  crystals  lose  their 
blue  colour,  and  become  white,  while  at  the  same  time  they  lose 
their  crystalline  nature  and  are  changed  to  a  powder.  Healing 
the  crystals  has  driven  off  their  water  of  crystallization,  and  conse- 
quently they  no  longer  retain  their  shape  as  crystals. 

In  the  case  of  the  copper  sulphate  it  is  easy  to  see  when 
the  water  of  crystallization  is  removed,  because  the  salt  itself 
(the  anhydrous  salt,  as  we  call  it)  is  nearly  white,  while  the 
crystallized  salt  is  deep  blue.  If  a  quantity  of  the  white  de- 
hydrated copper  sulphate  be  moistened  with  water,  it  is  at 
once  re-hydrated,  that  is,  it  takes  up  water  of  crystallization 


Water  of  Crystallization.  83 

again  and  consequently  turns  blue.  Many  other  salts  change 
colour  when  they  lose  either  some  or  all  of  their  water  of 
crystallization. 

Experiment  84. — Dissolve  a  crystal  of  cobalt  chloride  in  water. 
Notice  the  pink  colour  of  the  salt  and  also  of  the  solution.  With 
a  brush,  or  a  clean  pen,  write  on  paper,  using  this  solution  instead 
of  ink.  The  writing  will  be  invisible,  because  the  pink  colour  is 
so  faint.  Now  warm  the  paper  in  front  of  a  fire,  or  over  a  gas 
flame,  and  observe  that  the  writing  begins  to  show,  and  appears 
blue.  The  pink  salt  has  now  lost  water  of  crystallization,  and  turns 
blue.  If  the  paper  be  left  exposed  to  the  moisture  in  the  air,  or 
more  quickly  if  breathed  upon,  the  salt  will  re-hydrate  itself  and 
turn  pink  again. 

Some  salts  lose  their  water  of  crystallization  on  mere 
exposure  to  the  air.  Common  washing  soda  is  an  example. 
If  crystals  of  this  are  left  on  the  table  they  soon  lose  their 
clear  appearance  and  begin  to  fall  to  powder.  This  process 
is  called  efflorescence. 

Other  salts  do  just  the  opposite.  They  absorb  moisture 
from  the  air ;  sometimes  enough  to  cause  the  salt  to  liquify. 
Such  salts  are  said  to  deliquesce.  Substances  which  have  this 
power  are  very  useful  to  chemists  for  removing  moisture  from 
gases  which  are  required  to  be  dry. 

HYDROGEN  PEROXIDE. 

Besides  water,  hydrogen  and  oxygen  form  another  compound 
called  hydrogen  peroxide.  It  is  obtained  by  acting  on  either 
sodium  peroxide  or  barium  peroxide  with  dilute  sulphuric  acid. 
It  is  a  very  unstable  substance,  being  easily  converted  into  oxygen 
and  water.  The  readiness  with  which  it  parts  with  oxygen  enables 
it  to  oxidize  other  substances.  Thus,  if  it  be  added  to  lead  sulphide 
(a  black  compound),  it  converts  it  into  lead  sulphate  (a  white 
substance).  It  is  used  on  this  account  to  clean  old  oil  paintings 
which  have  become  black  by  the  white  lead  in  the  paint  being 
changed  to  lead  sulphide.  Hydrogen  peroxide  also  has  bleaching 
properties,  and  is  sometimes  employed  to  bleach  the  hair. 

EPITOME. 

Water  is  a  compound  of  oxygen  and  hydrogen.  This  is  proved 
synthetically  (i)  by  burning  hydrogen  in  oxygen,  and  collecting 


84  Wafer. 

and  testing  the  liquid  product  that  is  formed.  (2)  By  exploding 
oxygen  and  hydrogen  in  a  closed  vessel  (Cavendish's  method). 
(3)  By  reducing  certain  metallic  oxides,  such  as  copper  oxide,  in 
a  stream  of  hydrogen  (Duma's  method).  It  may  be  proved 
analytically  (i)  by  decomposing  water  by  certain  metals,  as  sodium, 
potassium,  magnesium,  or  iron:  whereby  oxides  of  the  metals 
are  formed  and  hydrogen  liberated.  (2)  By  heating  steam  with 
chlorine,  whereby  hydrogen  chloride  (hydrochloric  acid)  is  formed 
and  oxygen  set  free.  (3)  By  the  electrolytic  decomposition  of 
water,  whereby  both  oxygen  and  hydrogen  are  liberated. 

Both  by  synthesis  and  analysis,  we  learn  that  water  contains 
its  two  constituents  combined  in  the  proportion  of  one  part  by 
weight  of  hydrogen,  to  eight  parts  by  weight  of  oxygen ;  and  also 
that  the  proportion  by  volume  is  two  volumes  of  hydrogen  to  one 
volume  of  oxygen. 

When  two  volumes  of  hydrogen  and  one  volume  of  oxygen 
combine,  they  give  two  volumes  of  steam.  When  seen  in  mass, 
water  appears  bluish-green.  Its  point  of  maximum  density  is  4°  C. 
When  it  freezes  it  expands  ^  of  its  volume. 


CHAPTER   IX. 

WEIGHING    AND    MEASURING. 

As  soon  as  chemists  began  to  study  chemical  changes  from  a 
quantitative  point  of  view;  that  is  to  say,  as  soon  as  they 
realized  that  certain  relations  existed  between  the  quantities 
of  the  substances  which  take  part  in  chemical  changes,  it  then 
became  important  to  be  able  to  measure  and  to  weigh  with 
great  accuracy ;  and  as  chemistry  has  become  more  and  more 
an  exact  science,  the  various  methods  of  measuring  and  weighing 
have  developed  in  accuracy  and  delicacy. 

The  Metric  System. — The  measures  and  weights  used  in 
all  scientific  work  are  those  of  the  French  metric  system.  It 
is  called  the  metric  system  because  the  standard  unit  of  length 
is  the  metre.  This  corresponds  roughly  to  the  English  yard, 
being  39*37  inches,  and  to  this  measure  the  standards  ot 
capacity  and  of  weight  are  very  simply  related. 

(i)  Measures  of  Length. — The  metre  is  divided  into  tenths, 
hundredihs,  and  thousandths,  which  are  called  respectively, 
decimetres,  centimetres,  and  millimetres.  There  are  therefore 
10  millimetres  in  i  centimetre,  and  10  centimetres  in  i  deci- 
metre. 

The  relation  between  these  subdivisions  and  the  English 
inch  will  be  seen  by  Fig.  58,  which  shows  side  by  side  a  deci- 
metre scale,  divided  into  ten  centimetres,  and  each  of  these 
again  into  ten  millimetres  ;  and  a  four-inch  scale,  divided  into 
sixteenths.  Roughly  speaking,  2 '5  cm.,  or  25  mm.,  are  equal 
to  i  inch. 

For  the  most  part,  the  chemist  uses  only  one  measure  of 


86 


Weighing  and  Measuring. 


I 


length,  viz,  the  millimetre ;  he  records  a  length  as  20  mm.,  or 
760  mm.,  as  the  case  may  be,  instead  of  2  cm.,  or  7  dm.  6  cm. 
1000  metres,  or  i  kilometre,  is  a  little  over  half  a  mile 
(0*62  miles),  and  is  the  unit  employed  on  the  continent  of 
<>>  Europe  for  measuring  distances.     Instead  of 

milestones,  kilometre  stones  are  employed  upon 
the  roads. 

(2)  Measures  of  Capacity,  or  Volume. — The 
.     space  occupied  by  a  cube  whose  sides  measure 

one  centimetre  (see  dotted  cube,  Fig.  58); 
that  is  to  say,  the  volume  of  one  cubic  centi- 
metre is  taken  as  the  unit.  As  there  are  10 
centimetres  in  a  decimetre,  there  will  be 
10X10X10=  1000  cubic  centimetres  in  a 
cubic  decimetre. 

A  cubic  decimetre  is    called   a    litre;   a 
litre,   therefore,   contains    1000    cubic   centi- 
metres.    (The  litre  is  equal  to  about  if  pints, 
»  and  is  the  common  continental  unit  of  mea- 
ts sure  for  liquids.) 

(3)  Measures  of   Weight. — The  weight  of 
one  cubic  centimetre  (i   cc.)  of  pure  water 
(measured   at   a   temperature  when  water   is 
most  dense,  see  p.  79)  is  taken  as  the 
standard  unit  of  weight,  and  is  called  a 
gramme  (spelt  in  English,  gram).    Fig. 

59  represents  a  gram  weight  (in  brass), 
exact  size;   it  is  equal  to  about  15^ 
grains ;  and  31*1  grams  are  equal  to  i  oz.  Troy 
weight.     The  gram  is  subdivided  into  tenths, 
hundredths,  and  thousandths,  called  decigrams, 
centigrams,  and  milligrams  respectively. 

The  kilogram  (frequently  called  simply  a 
"kilo")  is  1000  grams.      It  is  equal  to  about 
2\  Ibs.,  and  is  the  common  continental  unit. 
Since  i  cc.  of  water  weighs  i  gram,  it  is  obvious  that  we 
at  once  know  the  weight  of  any  volume  of  water  expressed  in 
cubic  centimetres ;  or,  vice  versa,  we  know  the  volume  of  any 


10 


—  C\| 


— tai 


Thermometers.  87 

weight  of  water  given  in  grams.     For  instance,   1000  cc.  of 
water—  that  is,  i  litre  —  weighs  1000  grams,  or  i  kilo. 

Similarly,  if  we  know  the  specific  gravity  of  a  liquid  (that 
is,  how  many  times  lighter  or  heavier  than  water  it  is)  we  can 
easily  find  out  the  weight  in  grams  of  any  volume  of  it 
expressed  in  cubic  centimetres,  or  vice  versd.  For  example, 
suppose  we  have  200  cc.  of  a  liquid  whose  specific  gravity  is 
0-5,  and  we  wish  to  know  how  much  it  weighs  :  — 

Then,  as  i  :  0*5  :  :  200  :  x, 
x  =  200  X  0-5  =  100  grams. 

Or,  again,  the  specific  gravity  of  mercury  is  13*6,  what  will  be 
the  volume  of  408  grams  of  this  liquid  ? 

Then,  as  13*6  :  i  :  :  408  :  x, 


=        cc 
13*6 

Instruments  for  Weighing  and  Measuring.—  (i) 
Thermometers.  These  are  instruments  for  measuring  tem- 
perature. Their  use  depends  upon  the  fact  that  liquids  expand 
when  warmed,  and  contract  when  cooled.  The  liquid  most 
commonly  employed  is  either  mercury  or  alcohol.  Water 
would  not  be  suitable,  because  when  moderately  cooled  it 
freezes.1  The  graduations,  or  degrees,  upon  the  stem  of  a 
thermometer  are  purely  arbitrary  divisions.  The  instrument 
is  first  placed  in  a  vessel  filled  with  broken  ice,  and  the 
position  of  the  liquid  in  the  stem  is  marked  upon  the  glass. 
It  is  then  placed  in  steam,  from  water  which  is  kept  briskly 
boiling.  The  liquid  in  the  thermometer  at  once  expands  and 
rises  in  the  stem,  and  the  point  to  which  it  reaches  is  also 
marked  upon  the  glass.  These,  then,  are  the  two  fixed  points 
of  the  scale,  namely,  the  temperature  of  melting  ice  and  of 
steam  from  boiling  water.  The  space  upon  the  stem  between 
these  two  fixed  points  is  afterwards  divided  into  a  certain 

1  The  student  must  refer  to  text-books  on  physics  for  details  of  the 
method  of  making  and  graduating  thermometers. 


88 


Weighing  and  Measuring. 


number  of  equal  parts,  and  the  divisions  continued  both  above 
and  below.  It  is  quite  optional  how  many  equal  divisions  the 
space  between  the  two  fixed  points  shall  be  divided  into.  For 
instance,  we  might  decide  to  divide  it  into  ten  equal  parts, 
calling  the  lower  starting  point  o.  Then,  on  our  scale,  the 
melting  point  of  ice  would  be  o° ;  the  boiling  point  of  water 
would  be  10°;  blood  heat  would  be  37°;  and  the  ordinary 
temperature  of  a  warm  room  would  be  1-5°;  but  it  will  be 
R  c  F  obvious  that  it  would  be  best  to 
adopt  one  recognized  scale.  Unfor- 
tunately there  are  three  scales  in 
common  use.  In  one,  the  space  be- 
tween the  fixed  points  is  divided 
into  80  equal  parts,  the  lower  starting 
point  being  the  zero.  In  the  second, 
it  is  divided  into  100  equal  parts,  the 
gzero  being  the  same.  In  the  third, 
lit  is  divided  into  108  equal  parts,  and 
^the  zero  is  placed  32  divisions  below 
§  the  lower  fixed  point. 

The  first  of  these,  called  the 
Reaumur  thermometer  (R.  Fig.  60), 
is  the  one  in  common  use  on  the  con- 
tinent of  Europe.  The  third  is  known 
as  the  Fahrenheit  thermometer  (F.), 
and  is  commonly  used  in  England  for 
ordinary  purposes.  As  measured  by 
this  scale,  the  melting  point  of  ice 
(or  the  freezing  point  of  water)  is 
32°,  and  the  boiling  point  of  water 
The  other  is  the  Celsius,  or  more  commonly  known 
as  the  Centigrade  (C.)  thermometer,  and  is  always  used  for 
scientific  purposes. 

The  relation  in  which  these  three  scales  stand  to  each 
other  will  be  evident  from  the  figure.  It  is  obvious  that  the 
divisions  on  the  C.  scale  are  nearly  twice  as  long  as  those  of 
the  F.  scale,  100  C.  divisions  being  equal  to  180  F.  degrees, 
or  i°C.  •=  i  8°  F. 


FIG.  60. 


IS  212 


The  Barometer.  89 

In  order,  therefore,  to  translate  temperatures  given  upon 
the  centigrade  scale  to  degrees  Fahrenheit,  it  is  only  necessary 
to  multiply  by  r8,  and  add  32,  because  the  zero  of  the  latter 
scale  is  placed  32  divisions  below  the  point  at  which  the  centi- 
grade scale  starts.  And  vice  versd,  to  convert  degrees  F.  into 
degrees  C,  we  must  first  subtract  32  and  then  divide  by  i'8  : — 

(»°  C.  x  1-8)  +  32  =  °F.  ;  and  ?L°A.- 3?  =  °C. 

I'o 

Example  (i).1 — The  temperature  of  a  bath  is  95°  C. 
What  would  this  be  on  the  F.  scale  ? 

95  x  r8  =  171-6  ;  171-6  +  32  =  203-6. 

/.  95°  C.  =  203-6°  F. 

Example  (2). — The  temperature  recorded  on  a  hot  summer 
day  was  95°  F.  What  is  this  on  the  C.  scale  ? 

_  A3  = 

.'.  95°  F.  =  35°  C. 

(2)  The  Barometer  is  an  instrument  for  measuring  the 
pressure  of  the  atmosphere.  In  its  simplest  form  it  consists 
of  a  long  straight  glass  tube  closed  at  one  end,  which  has 
been  filled  completely  with  mercury,  and  inverted  with  its 
open  end  in  a  dish  of  mercury. 

Experiment  8$. — Take  a  stout  glass  tube  about  a  metre  long 
and  seal  up  one  end  in  the  blow-pipe  (see  p.  37).  Pour  mercury 
into  the  tube  until  it  is  nearly  full,  firmly  close  the  open  end  with 
the  thumb,  and  tip  the  tube  so  as  to  allow  a  large  air  bubble  to 
travel  to  the  other  end.  This  will  sweep  out  all  the  small  bubbles 
which  were  sticking  to  the  glass.  Now  completely  fill  up  the  tube 
with  mercury,  once  more  close  it  with  the  thumb  and  invert  it,  and 
lower  the  end  into  a  dish  or  trough  containing  mercury  and  care- 
fully withdraw  the  thumb.  Notice  that  the  mercury  falls  in  the 
tube  to  a  certain  point,  and  there  remains  stationary  (a,  Fig.  61). 

1  The  student  should  set  himself  a  few  such  examples  to  work  out. 
For  instance,  he  may  verify  the  various  temperatures  which  are  shown  to 
correspond  on  the  scales  in  Fig.  60. 


go  Weighing  and  Measuring. 

Now,  why  does  the  mercury  fall  in  the  tube  at  all  ?  and 
why,  if  it  falls  at  all,  does  it  not  drop  down  altogether?  In 
order  to  answer  these  questions  let  us  first  inquire  what  there 
is  in  the  clear  space  above  the  mercury.  Can  it  be  air? 
Unless  a  little  air  was  accidentally  allowed  to  get  in  when  the 
tube  was  inverted  in  the  mercury  dish,  it  can  hardly  be  air, 
because  the  tube  was  entirely  filled  with  mercury,  and  the 
thumb  was  not  removed  until  after  the  mouth  of  the  tube  was 
actually  under  the  surface  of  the  liquid.  Try  the  following 
experiment : — 

Experiment  86. — Slowly  tilt  the  tube,  and  notice  that  the 
mercury  goes  up  nearer  and  nearer  to  the  top,  until  at  last  the 
tube  is  as  completely  full  of  mercury  as  it  was  before  it  was  put 
into  the  trough.  This  proves  that  there  is  no  air  or  any  other  gas 
in  the  tube.  Raise  it  once  more  to  the  perpendicular  position  and 
the  mercury  again  falls  in  the  tube. 

This  experiment  shows  that  there  is  nothing  in  the  space 
above  the  mercury,  but  that  it  is  a  vacuum.1  This  being  the 
case,  it  follows  that  if  we  were  to  connect  one  end  of  a  long 
tube  to  the  most  perfect  air-pump,  and  dip  the  lower  end  in 
mercury,  it  would  be  impossible  to  pump  the  mercury  up 
the  tube  to  a  greater  height  than  it  stands  in  the  tube  in  the 
barometer,  because  in  that  tube  there  is  a  perfect  vacuum  in 
the  space  above  the  mercury. 

Experiment  87. — Measure  the  height  of  the  column  of  mercury 
in  the  tube,  measuring  from  the  surface  of  the  liquid  in  the  dish 
to  the  top  of  the  arched  surface  (called  the  mamscus\  a,  Fig.  61, 
of  that  in  the  tube.  It  will  be  found  to  be  about  760  mm.  [If  no 
millimetre  measure  is  at  hand,  use  a  foot  rule,  which  will  give 
nearly  30  inches  as  the  height.] 

Now  stand  the  apparatus  against  a  wall,  and  make  a  mark 
opposite  the  top  of  the  mercury  in  the  tube,  and  rule  a  horizontal 
line  along  the  wall  at  that  height.  Then  gradually  tilt  the  tube 
into  the  positions  /,  /',  Fig.  61,  and  notice  that  all  the  time,  the  top 
of  the  mercury  keeps  level  with  the  horizontal  line,  that  is  to  say,  it 
keeps  at  exactly  the  same  actual  vertical  height. 

1  This  experiment  was  first  made  by  the  Italian  physicist  Torricelli, 
and  the  space  is  called  the  Torricellian  Vacuum. 


The  Barometer. 


The  mercury  is  kept  up  in  the  tube,  because  the  atmosphere 
pressing  down  upon  the  mercury  in  the  dish,  pushes  it  up 
until  the  weight  of  liquid  in  the  tube  exactly  balances  the 
pressure  of  the  atmosphere.  If  from  any  causes  the  pressure 
of  the  atmosphere  be- 
comes greater,  it  will 
push  the  mercury  higher 
up  in  the  tube ;  and,  in 
the  same  way,  if  the  at- 
mospheric pressure  be- 
comes less,  it  cannot  sup- 
port so  long  a  column  of 
mercury,  and  therefore 
the  liquid  sinks  a  little 
in  the  tube.  It  is  per- 
fectly easy  to  show  by 
experiment,  that  it  is  the 
pressure  of  the  air  upon 
the  liquid  in  the  dish 
which  keeps  the  mercury 
up  in  the  tube.  If,  for 
instance,  we  surround 
the  apparatus  with  a  tall  glass  shade,  and  with  an  air-pump 
gradually  pump  the  air  out  of  the  shade,  we  should  notice  that 
the  mercury  would  gradually  sink  lower  and  lower  in  the  tube. 
By  the  time  we  had  pumped  out  half  the  air  (that  is,  reduced 
the  pressure  of  the  air  on  the  mercury  in  the  dish  to  one  half), 
we  should  find  that  the  mercury  was  now  standing  at  only  half 
its  original  height  in  the  tube,  namely,  380  mm.  instead  of  760. 

Again,  if  we  were  to  break  open  the  top  of  the  long  tube, 
the  mercury  would  of  course  instantly  sink  down  to  the  same 
level  as  that  in  the  dish,  because  the  atmospheric  pressure 
would  then  be  the  same  on  the  liquid  inside  and  outside  the 
tube. 

The  pressure  of  the  atmosphere  is  different  at  different 
parts  of  the  earth's  surface ;  and  is  also  liable  to  vary  in  the 
same  locality,  from  hour  to  hour.  Consequently  the  height 
of  the  column  of  mercury  which  it  is  able  to  support  also 


FIG.  61. 


92  Weighing  and  Measuring. 

varies;  therefore  the  barometer  enables  us  to  tell  at  any 
moment  what  is  the  actual  atmospheric  pressure  at  the  time. 
What  is  called  the  standard  or  normal  pressure  is  that  which 
is  able  to  support  a  column  of  mercury  760  mm.  high. 

(3)  The  Balance. — In  order  to  weigh  any  substance 
accurately,  or  to  detect  minute  differences  in  the  weights  of 
different  things,  it  is  necessary  to  employ  a  delicate  balance 
and  exact  weights.  For  instance,  a  few  grains  of  sand 
sprinkled  upon  the  scale  of  an  ordinary  kitchen  balance  will 


FIG.  62. 

make  absolutely  no  difference,  whereas  a  single  particle  of 
sand  placed  upon  one  pan  of  a  chemical  balance  would  com- 
pletely weigh  it  down.  The  chemical  balance  is  usually 
enclosed  in  a  glass  case,  partly  to  protect  it  from  dust  and 
dirt,  and  partly  in  order  that  it  may  not  be  exposed  to  the 
slightest  draught  when  being  used.  Such  a  balance  will 
readily  weigh  to  a  fraction  of  a  milligram. 

For  the  purposes  of  the  elementary  student,  an  instru- 
ment of  the  extremest  delicacy  is  not  necessary  or  desirable. 
Fig  62  shows  a  balance  suitable  for  his  requirements. 

By  turning  the  handle  or  lever  H,  the  beam  is  liberated 
from  its  support,  and  is  then  free  to  swing.  This  balance 


The  Balance.  93 

will  turn  with  2  milligrams,  and  is  able  to  carry  as  much  as 
100  grams.  It  must  not  be  used  for  heavier  weights  than 
this. 

The  set  of  weights  consists  of  the  following,  50,  20,  10, 
TO,  5,  2,  i,  i,  i,  grams,  making  up  100  grams;  and  0-5,  0-2, 

O'l,  O'l,  0*05,  O'02,   O'OI,  O'OI. 

In  order  to  make  a  weighing,  we  proceed  as  follows — 

Experiment  88. — Place  the  object  to  be  weighed,  say  a  clean 
porcelain  crucible  with  its  lid,  upon  the  left  pan,  and  put  on  the 
other  pan  a  weight  which  is  roughly  judged  to  be  equal  to  it,  say 
the  20-gram  weight.  Use  the  forceps  (not  the  fingers)  to  lift  the 
weights,  and  place  them  gently  upon  the  scale  pan.  Now  release 
the  beam,  by  means  of  the  lever,  and  observe  which  scale  pan 
falls  ;  we  will  suppose  the  2o-gram  weight  is  too  much.  Raise 
the  beam  again,  lift  off  the  weight,  return  it  to  its  place  in  the  box, 
and  put  the  lo-gram  weight  on  the  pan.  Suppose  this  is  too  little, 
put  on  the  5-gram  in  addition.  The  weights  must  never  be 
removed  or  added  while  the  beam  is  free,  but  only  when  it  is 
supported  in  the  rest.  If  15  grams  is  too  much,  remove  the  5, 
return  it  to  its  place,  and  put  on  the  2-gram  :  if  this  is  too  little 
put  on  i  gram,  and  if  this  is  too  much  then  the  crucible  weighs 
more  than  12  but  less  than  13  grams.  Now  add  0*5  gram  ;  suppose 
this  to  be  too  heavy,  remove  it  and  try  o'2.  If  this  is  too  little  add 
0*1,  if  now  too  much  remove  the  OT  and  put  on  0-05.  Suppose  this 
to  be  the  exact  weight,  then,  as  the  beam  oscillates,  the  pointer  will 
swing  to  the  same  distance  upon  the  scale  on  both  sides,  and  the 
weight  of  the  crucible  is  12*25  grams.  [NOTE. —  When  not  in  use, 
the  balance  and  the  weights  should  not  be  left  exposed  to  the 
laboratory  atmosphere,  but  shoidd  be  either  covered  ivith  a  glass 
shade,  or  put  away  into  a  clipboard^ 

(4)  For  measuring  liquids,  graduated  glass  vessels  are 
employed.  For  moderate  volumes,  a  graduated  250  cc. 
cylinder  may  be  used  (Fig.  63).  While  for  small  quantities  a 
burette  is  more  useful  (Fig.  64). 

When  small  definite  volumes  are  required,  say  exactly 
10  cc.  or  25  cc.,  a  pipette  is  employed.  This  is  dipped  into 
the  liquid,  which  is  then  sucked  up  the  tube,  nearly  to  the 
top,  by  the  mouth,  and  the  top  quickly  closed  with  the  finger 
(Fig  65).  By  cautiously  releasing  the  finger,  the  liquid  is 


94  Weighing-  and  Measuring. 

allowed  to  drop  out  until  it  sinks  exactly  to  the  mark  on  the 
narrow  stem.  When  larger  definite  volumes  are  required, 
flasks  holding  i,  £,  or  i  litre  are  used. 

Experiment  89.— Select  a  flask  with  rather  a  narrow  neck,  and 
of  such  a  size  that  when  500  cc.  of  water  have  been  measured  into 
it  (by  twice  filling  the  250  cc.  cylinder)  the  water  stands  in  the 
neck  of  the  flask.  Now  gum  a  label  round  the  neck  in  such  a 
position  that  the  top  edge  of  the  label  is  exactly  level  with  the 


TOM 


FIG.  63. 


FIG.  65. 


FIG.  64. 

water.  If  now,  by  means  of  a  file,  a  slight  scratch  be  made  upon 
the  glass,  just  where  the  upper  edge  of  the  label  is,  a  half  litre 
measuring  flask  will  have  been  made. 

The  volume  which  a  given  weight  of  a  liquid  occupies, 
however,  depends  upon  the  temperature  of  the  liquid. 

Experiment  90.— Fill  the  half  litre  flask  exactly  to  the  mark 
with  cold  water ;  place  the  flask  upon  a  piece  of  wire  gauze  upon 
a  tripod  and  gradually  warm  it  with  a  gas  flame.  Notice  that  the 
water  rises  in  the  neck  of  the  flask  some  distance  above  the  mark. 


Measuring  Gases.  95 

The  water  has  expanded^  or  increased  in  bulk.  Now,  by  means  of 
a  pipette,  withdraw  some  of  the  water,  until  exactly  the  half  litre  is 
left,  and  let  the  flask  cool  again.  When  it  is  cold,  observe  that  it 
no  longer  contains  as  much  as  half  a  litre  of  water. 

This  shows  that,  if  we  measure  a  volume  of  liquid  at  a  low 
temperature,  we  have  a  greater  actual  weight  of  it  than  if  we 
measure  it  at  a  high  temperature,  and  therefore  we  ought  to 
measure  always  at  the  same  temperature,  or  else  make  a  proper 
allowance  for  the  expansion  or  contraction.  In  ordinary 
experiments,  when  extreme  accuracy  is  not  necessary,  and  the 
quantities  are  only  small,  measurements  made  at  the  ordinary 
temperature  of  the  room  are  regarded  as  being  made  at  the 
same  temperature,  and  no  correction  is  made.  The  wholesale 
spirit  dealer,  on  the  other  hand,  takes  careful  note  of  the 
temperature  at  which  his  spirit  is  measured,  else  if  he  were 
to  buy  large  volumes  in  hot  weather,  and  sell  it  in  cold 
weather,  he  would  be  a  considerable  loser  by  the  transaction. 

(5)  Measuring  Gases. — We  may  measure  the  volume 
of  a  gas  obtained  in  a  chemical  process,  by  collecting  it  over 
water  or  mercury  in  a  vessel  graduated  into  cubic  centimetres ; 
or  we  can  collect  it  in  an  urigraduated  vessel,  mark  the  volume 
by  gumming  a  label  upon  the  glass,  and  afterwards  seeing  what 
volume  of  water  has  to  be  poured  into  the  vessel  from  the 
250  cc.  measure,  to  fill  it  up  to  the  mark.  This  is  a  compara- 
tively rough  method  ;  in  more  exact  work  it  is  usual  to 
measure  gases  in  glass  tubes  over  mercury  by  means  of  .a 
millimetre  scale.  Vessels  used  for  measuring  gases  are  called 
eudiometers.  Gases  expand  when  warmed,  and  contract  on 
being  cooled  just  as  liquids  do,  but  to  a  much  greater  extent. 

Experiment  91. — Fit  a  cork  and  delivery  tube  into  a  dry  half- 
litre  flask,  and  arrange  it  as  shown  in  Fig.  66.  Gently  heat  the 
empty  flask  with  a  lamp,  and  notice  that  the  air  within  the  flask 
quickly  expands,  and  escaping  through  the  delivery  tube,  is  col- 
lected in  the  cylinder  over  water.  Observe  that  the  volume  of  the 
gas  which  is  driven  out  of  the  flask  is  much  greater  than  the  volume 
of  liquid  which  was  removed  from  the  flask  of  water  in  Exp.  90  by 
the  pipette.  Remove  the  flame,  and  notice  that,  as  the  air  in  the 


96  Weighing  and  Measuring. 

flask  cools,  and  therefore  contracts,  water  is  drawn  back  from  the 
trough  to  take  its  place. 

When  measuring  the  volume  of  gases,  therefore,  it  is  always 


FIG.  66. 


necessary  to  take  into  account  the  temperature  of  the  gas  at  the 
time,  and  then  to  calculate  what  the  observed  volume  would 
be  if  the  temperature  were  to  be  lowered  to  o°  C,  which  is 
taken  as  the  standard  or  normal  temperature. 


CHAPTER  X. 

SOME   GENERAL    PROPERTIES    OF    GASES. 

Relation  of  the  Volume  of  Gases  to  Heat.— Nearly  a 
century  ago  it  was  discovered  by  Charles  and  Gay-Lussac  that 
all  gases  expand  and  contract  to  the  same  extent  under  the  same 
changes  of  temperature,  provided  there  is  no  alteration  in  the 
pressure.  This  is  known  as  the  Law  of  Charles. 

For  every  degree  of  the  centigrade  scale  that  a  gas  is  heated, 
it  expands  -5-f  3-  part  of  the  volume  it  occupies  at  o°  C. 

Thus,  i  vol.  at  o°  C.  becomes  i  -f  ^  at  i°,  and  i  +  ^ 
at  2°,  and  so  on. 

Or  273  vols.  at  o   become  273  +  1,  or  274  at  i° 

273  +  2,  „  275  „  2° 

273  +  3,  „  276  „  3°,  etc. 

„    v     >°      ,    273  +  /,         „  f 

If  V°  stands  for  the  volume  at  o°  C.,  and  V  for  the  volume 
at  /  temperature,  then  we  get  the  simple  proportion — 

273:273  +  '::^:  V, 


273 

EXAMPLE  i. — Suppose  we  have  a  quantity  of  gas  which 
measures  320  cc.,  while  the  temperature  is  15°  C.  What  volume 
will  this  occupy  at  o°  ? 

Then  V°  =  32°  X273  =  303-3  cc. 
273  +  15 

EXAMPLE  2.— 250  cc.  of  gas  measured  at  -  10°  C.  What  will 
be  the  volume  at  the  normal  temperature  ? 

Then  V°  =  Hl°  x  273  =  2$    $ 
273  -  10 


Some  General  Properties  of  Gases. 


EXAMPLE  3.  —  A  quantity  of  gas  occupies  100  cc.  when  measured 
°.     What  will  it  measure  when  heated  to  30°  ? 


Then  V  =  _      =  „       cc. 

273 

EXAMPLE  4.—  560  volumes  of  gas  measured  at  10°,  are  heated 
to  20°.     How  many  volumes  will  they  then  occupy  ? 

Then,  273  +  10  :  273  +  20  :  :  560  :  V 

or,  V  =  56o  x  (273  +.20)  =  5?6  volumes 
273  +  10 

Relation  of  the  Volume  of  Gases  to  Pressure.  —  If 

we  squeeze  a  solid  or  a  liquid,  we  observe  no  change  in  their 
volumes  ;  but  if  we  put  pressure  upon  a 
gas,  it  at  once  becomes  greatly  reduced 
in  volume.  Robert  Boyle  (1661)  dis- 
covered that  the  volume  occupied  by  a 
given  weight  of  any  gas  is  inversely  as  the 
pressure.  This  is  known  as  Boyle's  Law. 
It  means  that  if  we  double  the  pressure 
on  a  gas  we  reduce  the  volume  to  one- 
half;  while  if  we  diminish  the  pressure 
to  one-half,  we  increase  the  volume  to 
double.  It  is  easy  to  put  this  law  to  the 
test.  Suppose  a  quantity  of  air  is  en- 
closed in  the  tube  T  (Fig.  67),  which  is 
connected  by  means  of  a  thick  india- 
rubber  tube  to  a  small  reservoir  of  mer- 
cury, capable  of  being  hoisted  up  and 
down.  Let  us  first  place  the  reservoir 
in  such  a  position  that  the  mercury  in 
it  is  exactly  level  with  that  in  the  tube. 
Under  these  circumstances  the  enclosed 
air  is  under  the  ordinary  atmospheric 
call  it  i  vol.,  and  mark  it  upon  the  tube. 

Next  we  will  place  a  mark  at  an  equal  distance  below,  and 

another  at  one-half  the  distance  above  it 

Now  suppose  we  raise  the  mercury  reservoir,  we  shall  see 

the  volume  of  the  gas  diminishing  until  presently  the  mercury 


FIG.  67. 

pressure.     Let 


us 


The  Relation  of  Volume  to  Pressure.  99 

stands  at  the  i-mark.  When  it  reaches  this  point  we  shall 
find,  on  measuring,  that  the  height  of  the  mercury  in  the 
reservoir  above  that  in  the  tube  is  about  760  mm.  We 
have  already  learnt  that  760  m;n.  of  mercury  represents  a 
pressure  equal  to  that  of  the  atmosphere,  hence  we  have 
subjected  the  gas  within  the  tube  to  a  pressure  of  an  additional 
atmosphere ;  that  is,  there  is  now  twice  the  pressure  upon  it 
that  there  was  at  first,  and  its  volume  is  reduced  to  one-half. 

In  the  same  way,  if  the  reservoir  be  lowered  down  until  the 
gas  in  the  tube  has  expanded  to  twice  its  original  vol.,  that 
is.  down  to  the  2  mark,  we  should  find,  on  measuring,  that  the 
mercury  in  the  reservoir  was  standing  380  mm.  below  that  in 
the  tube.  Instead  of  being  under  the  ordinary  atmospheric 
pressure,  the  gas  is  now  under  reduced  pressure.  It  is  sub- 
jected to  a  pressure  of  760  mm.  —  380  mm.  =  380  mm.  That 
is,  to  a  pressure  of  only  half  the  ordinary  atmospheric  pressure, 
and  we  see,  as  demanded  by  Boyle's  law,  that  its  volume  is 
doubled. 

Experiment  92. — Take  a  glass  tube  about  half  a  metre  long,  and 
closed  at  one  end,  and  completely  fill  it  with  mercury.  Now  pour 
out  this  mercury  into  a  cc.  measure,  and  in  this  way  ascertain  the 
exact  capacity  of  the  tube.  Then  pour  back  into  the  tube  one-half 
the  mercury,  and  mark  upon  the  tube  where  it  comes  to.  This 
mark  will  therefore  indicate  exactly  half  the  capacity  of  the  tube. 
Now  close  the  tube  with  the  thumb  and  invert  it  in  a  dish  of 
mercury,  and  stand  it  in  a  vertical  position.  Notice  that  the 
mercury  does  not  stand  at  the  mark,  but  at  a  point  some  distance 
below  it.  Why  is  this,  for  the  tube  when  inverted  was  half  filled 
with  air  and  half  with  mercury  ?  The  reason  is  because  the  gas  is 
under  reduced  pressure,  and  has  therefore  expanded.  It  is  under  a 
pressure  equal  to  that  of  the  atmosphere  minus  that  of  the  column 
of  mercury  in  the  tube.  We  must,  therefore,  first  ascertain  the  actual 
height  of  the  barometer  at  the  time,  and  then  carefully  measure  the 
height  of  the  column  of  mercury,  from  the  surface  of  that  in  the  dish 
to  the  top  of  that  in  the  tube.  Suppose  the  barometer  happens  to 
be  low,  say  740  mm.,  and  the  length  of  the  column  in  the  tube  to 
be  200  mm.,  then  740  -  200  =  540  mm.  is  the  pressure  to  which  the 
gas  is  exposed. 

Experiment  93. — Take  the  tube  used  above,  fill  it  with  water 
up  to  the  mark  and  invert  it  in  a  dish  of  water.  Notice  where  the 


ioo  Some  General  Properties  of  Gases. 

water  stands  ;  it  only  falls  a  very  little  way  below  the  mark.  Why; 
is  this  ?  It  is  because,  water  being  so  much  lighter  than  mercury 
(13^  times  lighter),  the  short  column  of  it  in  the  tube  scarcely 
reduces  the  pressure  on  the  gas  at  all. 

Seeing  that  the  volume  of  a  gas  is  so  closely  dependent 
upon  the  pressure  to  which  the  gas  is  exposed,  it  will  be 
evident  that  if  we  wish  to  compare  volumes  of  gases,  we  must 
know  what  they  measure  under  the  same  conditions  of  pres- 
sure. Now  it  is  not  often  possible,  and  is  seldom  convenient, 
to  actually  put  gases  under  one  regular  pressure  before  mea- 
suring their  volumes,  so  the  plan  always  adopted  is  to  measure 
the  volume  of  the  gas  at  the  particular  pressure  it  happens  to 
be  under,  and  then  to  calculate  what  volume  it  would  occupy 
at  a  pressure  of  760  mm.,  this  being  the  standard  or  normal 
pressure.  This  is  called  correcting  for  pressure,  and  the  calcu- 
lation is  a  simple  proportion,  based  on  Boyle's  law. 

For  example,  ioo  cc.  of  a  gas  measured  at  380  mm.,  what 
volume  would  it  occupy  at  760  mm.  ? 

Then,  as  760  :  380  : :  ioo  :  x 
=  ioo  X  380  =        ca 
760 

Again,  a  quantity  of  gas  measured  500  volumes  at  780  mm. 
What  will  it  occupy  at  the  standard  pressure  ? 

As  760  :  780  : :  500  :  x 
*  =  552*-Z??  =  5 13  volumes. 

700 

In  practice,  the  corrections  for  both  temperature  (p.  97) 
and  pressure  are  usually  made  together.  For  example,1  a 
sample  of  gas  measured  360  cc,  its  temperature  was  15°  C, 
and  it  was  at  atmospheric  pressure  ;  but  the  barometer  at  the 
time  was  standing  at  750  mm. 

Then  3g°,  x  273  _  tne  temperature  correction  alone, 
273+15 

and  ^  °       *5°  _  tne  pressure  correction  alone. 
760 

1  The  student  should  make  himself  perfectly  familiar  with  this  method 
of  reducing  volumes  of  gas  to  the  normal  temperature  and  pressure,  by 
working  out  a  number  of  examples. 


The  Crith.  101 

Putting  the  two  together,  we  get  — 

360  x  273  x  750  _.  73710000  =  --6.7  cc 

(273  +  15)  X  760          218880 


Therefore,  360  cc.  of  gas  at  i5°"G  wJ  750  mite  '^'336*7  cc. 
at  the  normal  temperature  and  'ptessure.  v  ' 

The  Crith.—  One  litre  'oV  hydiogf.n  gas,  sri^asiJr6d<'  at  o°  C. 
and  760  mm.,  weighs  0*0896  gianib.  This  is  an  'ex'ttemely  im- 
portant figure  to  remember,  for  by  means  of  it  we  can  calculate 
the  weight  of  any  volume  of  any  gas,  so  long  as  we  know  the 
density  of  that  gas;  that  is,  how  many  times  heavier  than 
hydrogen  it  is.  For  instance,  if  we  learnt  that  oxygen  was 
sixteen  times  as  heavy  a  gas  as  hydrogen,  bulk  for  bulk,  then 
i  litre  of  oxygen  will,  of  course,  weigh  sixteen  times  as  much  as 
i  litre  of  hydrogen;  that  is,  0-0896  x  16,  or  1*42  grams,  at 
normal  temperature  and  pressure  (N.T.P.).  This  number, 
0*0896,  is  so  important  that  a  name  has  been  given  to  it  (just 
as  in  mathematics  the  name  of  the  Greek  letter  TT  is  used  to 
denote  the  number  3-1416,  the  ratio  between  the  diameter  and 
the  circumference  of  a  circle).  The  name  adopted  for  the 
weight  of  a  litre  of  hydrogen  measured  at  N.T.P.  is  the  Greek 
word  signifying  a  barleycorn,  crith.  It  is  used  symbolically 
for  a  little  weight,  and  does  not  mean  that  a  litre  of  hydrogen 
weighs  as  much  as  a  barleycorn.  We  say  that  a  litre  of 
oxygen  weighs  16  criths  ;  that  is,  simply  sixteen  times  ofo896. 
Another  important  number  to  be  remembered  in  this  connec- 
tion is  the  volume  of  i  gram  of  hydrogen,  measured  at  N.T.P. 
i  gram  of  hydrogen  measures  11*165  litres  at  N.T.P.  From 
this  number,  just  as  from  the  crith^  we  are  able  to  calculate 
the  weight  of  any  volume  of  any  gas  whose  density  is  known. 
Thus,  if  oxygen  is  sixteen  times  as  heavy  as  hydrogen,  then, 
since  i  gram  of  hydrogen  measures  11*165  litres,  obviously, 
16  grams  of  oxygen  will  measure  in65  litres.  In  other 
words,  11*165  litres  is  the  volume  which  will  be  occupied  by 
the  same  number  of  grams  of  a  gas,  as  expresses  the  density 
of  that  gas.  Suppose  we  have  three  gases  whose  densities  are 
respectively  14,  i8±,  and  22  (that  is,  one  is  14  times,  one  i8£ 
times,  and  the  other  22  times  as  heavy  as  hydrogen),  then  14 


IO2  Some  General  Properties  of  Gases. 

grams  of  one,  i8±  grams  of  the  second,  and  22  grams  of  the 
third  will  measure  11*165  litres. 

The  liquefaction  of  gases. — We  have  learnt  (p.  97)  that  when 
a  gas  is  heated  from  o°  to  i°  it  expands  by  o^;j  of  its  bulk  ;  if  it  be 
cooled  frorrup^tO  —  i°'ij: ,&*##«##?  '^}5  of  its  volume;  from  o°  to  —  2, 

-  3,  -  4,  etc".,  'ft  contracts' ^f  j,  ^>f  g,  and  %$s  respectively.     Now,  if 
this  law."hal(ls;'gp*dd., 'however,  m^ch  we  cool  a  gas,  it  would  follow 
that  if  4a"*qua*ntpfy'  Of*  gas:  be'  cooled  'from  o°  to   -273°,  it  would 
contract  f  £|  of  its  volume,  that  is,  it  would  occupy  no  volume  at  all. 

Now,  a  temperature  as  low  as  -  273°  has  never  yet  been  reached 
(this  temperature  is  sometimes  called  the  absolute  zero).  The 
lowest  degree  of  cold  which  has  been  obtained  as  yet,  is  about 
-220°,  but  before  this  point  is  reached  all  the  known  gases  except 
one,  and  that  one  is  hydrogen,  pass  into  the  liquid  condition.  Just 
as  steam,  when  cooled,  changes  from  the  gaseous  state  to  the  liquid, 
so  all  other  gases,  when  sufficiently  strongly  cooled,  change  from  the 
gaseous  to  the  liquid  condition.  Some  gases  require  very  little 
cooling  to  make  them  do  this,  while  others  require  to  be  exposed 
to  the  lowest  possible  temperature  in  order  to  make  them  change 
their  state.  Among  the  latter  class  are  oxygen  and  nitrogen,  and 
it  is  because  it  is  only  recently  that  chemists  have  been  able  to 
obtain  the  necessary  degree  of  cold,  that  these  gases  have  only  of 
late  years  been  obtained  in  the  liquid  state.  For  example,  oxygen 
requires  to  be  cooled  down  to  the  extremely  low  temperature  of 

-  181°  to  cause  it  to  pass  from  gas  to  liquid. 

There  is  little  doubt  but  that  before  long  a  sufficiently  low 
temperature  will  have  been  reached  in  order  to  produce  liquid 
hydrogen. 

As  a  gas  gets  near  to  the  temperature  at  which  it  turns  into  a 
liquid,  it  begins  to  depart  from  the  law  of  Charles  and  Gay-Lussac, 
and,  of  course,  when  it  liquefies,  and  therefore  ceases  to  be  a  gas,  it 
is  no  longer  subject  to  the  laws  which  govern  gases. 

Again,  it  is  found  that  when  gases  are  subjected  to  pressure, 
they  sooner  or  later  begin  to  depart  from  Boyle's  law,  and  finally  to 
change  their  state  from  gases  to  liquids.  A  most  remarkable  point, 
however,  about  the  effect  of  pressure  in  causing  the  liquefaction  of 
a  gas,  is,  that  the  gas  must  be  below  a  certain  temperature.  If  it  be 
above  this  temperature,  no  amount  of  pressure  will  squeeze  it  into 
the  liquid  state.  This  particular  temperature  is  different  for  each 
gas,  and  is  called  the  critical  temperature  of  the  gas.  For  example, 
the  critical  temperature  of  chlorine  is  141°.  This  means  that  if 
chlorine  is  heated  above  this  point,  no  amount  of  pressure  will 


Liquefaction  of  Gases.  103 

make  it  pass  into  the  liquid  state  ;  but  at  all  ordinary  temperatures 
of  the  air,  chlorine  is  far  below  its  critical  temperature,  therefore  it 
can  be  easily  condensed  to  the  liquid  state  by  mere  pressure. 

Again,  the  critical  temperature  of  ethylene  (see  p.  184)  is  about 
10°,  that  is,  just  a  trifle  below  the  ordinary  temperature  of  a  room  ; 
hence,  in  order  to  compress  this  gas  into  the  liquid  state,  it  must 
be  slightly  cooled,  so  as  to  bring  its  temperature  below  its  critical 
point. 

In  the  case  of  some  gases  the  critical  temperature  is  extremely 
low,  thus  in  the  case  of  oxygen  it  is  — 118*8°.  Therefore,  in  order 
to  compress  oxygen  into  a  liquid,  it  is  absolutely  necessary  to  cool 
it  down  to  this  intense  degree  of  cold  before  liquefaction  will  take 
place.  At  -  u8'8°  oxygen  requires  a  pressure  of  fifty  atmospheres 
to  liquefy  it.  The  more  it  is  cooled  below  this,  the  less  pressure  is 
needed  to  liquefy  it,  until  at  -  181°  it  passes  into  the  liquid  condition 
at  the  ordinary  atmospheric  pressure. 

The  critical  temperature  of  hydrogen  is  lower  still,  and  up  to 
the  present  no  artificial  cold  has  been  obtained  low  enough  to  cool 
hydrogen  in  quantity  down  to  this  point  ;  hence  hydrogen  has  not 
yet  been  obtained  as  a  coherent  liquid,  although  by  special  devices 
momentary  indications  of  liquidity  have  been  observed,  when  the 
hydrogen  appeared  as  a  froth  or  spray. 


CHAPTER  XL 

SIMPLE   QUANTITATIVE    MANIPULATIONS.1 

EVERY  one  knows  that  if  we  mix  two  substances  together,  say 
salt  and  sugar,  we  can  have  the  two  ingredients  present  in  any 
proportion  we  please.  We  could,  for  instance,  make  a  mixture 
containing  so  large  a  proportion  of  salt,  that  the  sugar  present 
could  not  be  tasted ;  or  the  sugar  might  so  preponderate  that 
we  could  not  taste  the  salt  in  the  mixture.  Now  a  most 
important  question  arises,  viz.  Can  substances  enter  into 
chemical  combination  in  this  manner  ?  That  is  to  say,  when 
two  elements  unite,  can  they  do  so  in  any  proportion  ?  does 
the  composition  of  the  compound  they  produce  vary  ?  If  two 
substances  can  combine  together  in  any  proportion,  just  as 
they  can  be  mixed  together,  then  obviously  the  composition 
of  the  compound  produced  will  depend  upon  the  quantities  of 
the  ingredients  that  were  employed  to  produce  it.  If,  on  the 
other  hand,  it  should  be  that  there  is  some  fixed  proportion  in 
which  elements  combine,  that  if  they  unite  it  must  be  in  some 
particular  proportions,  or  not  at  all,  then  it  would  follow  that 
any  given  compound  would  always  have  absolutely  the  same 
composition.  This  is  a  point  which  it  is  of  the  very  greatest 
importance  to  settle,  and  in  order  to  do  so  we  must  make 
experiments  upon  weighed  or  measured  quantities  of  sub- 
stances, and  carefully  weigh  or  measure  the  resulting  products. 

1  The  quantitative  experiments  have  been  arranged  as  far  as  possible 
so  as  to  make  use  only  of  such  knowledge  of  chemical  facts  as  the  student 
has  already  gained  from  the  earlier  chapters.  The  teacher  will  do  well  to 
substitute  or  a<ld  others  from  time  to  time,  It  is  well  that  the  student 
should  check  his  own  results  by  doing  duplicate  experiments. 


The  Combustion  of  Magnesium.  105 

In  other  words,  we  must   make  quantitative  experiments,  as 
distinguished  from  qualitative. 

We  can  make  our  quantitative  investigations  in  two  ways  : 
either  synthetically  or  analytically.  We  can  cause  weighed 
quantities  of  materials  to  enter  into  chemical  union,  and  weigh 
the  products ;  or  we  can  decompose  known  quantities  of 
certain  compounds,  and  weigh  or  measure  the  resulting 
substances. 

Experiment  94. — Trie  combination  of  magnesium  with 
oxygen.  We  have  learnt  that  when  substances  burn  in  the  air, 
they  are  combining  with  oxygen  to  produce  oxides.  Let  us  now 
burn  a  weighed  quantity  of  magnesium  and  weigh  the  resulting 
oxide.  Take  a  porcelain  crucible  with  its  cover,  and  counterpoise  it 
on  the  balance  (p.  92).  Instead  of  using  the  weights,  place  a  pill- 
box in  the  opposite  pan,  and  put  into  it  some  fine  shot,  until  the 
crucible  is  exactly  balanced.  This  little  box  of  shot  will  now  be 
the  "  tare  "  for  the  empty  crucible. 

Next  weigh  out  into  the  crucible  from  half  to  three-quarters  of 
a  gram  of  magnesium  ribbon,  which  has  been  previously  scraped 
or  rubbed  down  with  sandpaper 
to  remove  any  oxide  from  it,  and 
cut  into  short  pieces.  Weigh  it 
carefully,  and  note  the  weight.1 

Now  place  the  crucible  upon 
a  pipe-clay  triangle  supported  on 
a  tripod  or  the  ring  of  a  retort 
stand,  and  heat  it  (Fig.  68)  ;  ap- 
plying the  heat  first  gently  and 
then  more  strongly.  From  time 
to  time  slightly  raise  the  cover  by  FIG  6g 

means   of  tongs,  to   observe  the 

progress  of  the  combustion,  but   not  so   as  to  allow  any  of  the 
oxide  to  be  carried  away. 

When  the  whole  of  the  magnesium  has  burnt,  allow  the  crucible 
to  cool.  Then  place  it  on  the  balance,  put  the  "tare"  on  the 
opposite  scale,  and  weigh  the  contents  of  the  crucible.  The  first 

1  Instead  of  using  a  tare,  the  crucible  may  be  weighed,  first  empty, 
and  again  with  the  metal  in  it;  then,  by  deducting  the  first  from  the 
second  weight,  the  weight  of  magnesium  is  found.  The  student  will  find 
it  quicker,  and  will  be  less  likely  to  introduce  errors,  if  he  uses  tares  as 
described. 


106  Simple  Quantitative  Manipulations. 

thing  observed  will  be  that  it  has  gained  in  weight.     This  we  have 
learned  to  expect  from  Exps.  72  and  73.     Carefully  note  the  weight. 
The  difference  between  the  weight  of  the  magnesium  employed, 
and  that  of  the  oxide  obtained,  is  the  weight  of  oxygen  from  the 
atmosphere  which  has  combined  with  the  magnesium. 
Calculate  the  proportions  in  the  following  way  : — 
Suppose  the  weight  of  magnesium  used  was  0*53  gram,  and  that 
of  the  magnesium  oxide  produced  to  have  been  0*88  gram  ;  then, 
subtracting  the  former  from  the  latter,  we  get— 

0-88 


0-35  =  weight  of  oxygen 

which  has  combined  with  0-53  grains  of  magnesium. 
Then,  as  0-35  :  0-53  :  :  i  :  x 


therefore  the  proportion  in  which  oxygen  and  magnesium  have 
combined  together  in  this  experiment  is  i  part  by  weight  of  oxygen 
with  i  '5  parts  of  magnesium. 

In  this  experiment  the  magnesium,  being  heated  in  the  air, 
is  surrounded  with  an  unlimited  supply  of  oxygen  ;  therefore, 
if  we  were  to  repeat  the  process  a  number  of  times,  each  time 
employing  a  different  weight  of  the  metal,  and  were  to  find 
that  in  every  case  the  magnesium  and  oxygen  combined  in 
exactly  the  same  proportions,  we  should  have  good  ground  for 
supposing  that  these  two  elements  could  only  unite  in  some 
definite  proportion. 

It  will  at  this  point  be  interesting  to  find  out  whether  the 
same  result  would  follow  if  the  magnesium  were  to  be  burnt  in 
pure  oxygen,  instead  of  in  the  air.  For  this  purpose  the  next 
experiment  should  be  made. 

Experiment  95.  —  Take  a  piece  of  combustion  tube  (p.  32)  with 
moderately  thin  walls,  about  18  cm.  (7  inches)  long,  and  slightly 
border  the  ends  (see  p.  36).  Into  one  end  fit  a  cork  with  a  short 
straight  glass  tube,  and  insert  into  the  other  a  short  plug  of  cotton 
wool,  as  shown  in  Figure  69.  Now  counterpoise  the  whole  as  in 


The  Combustion  of  Magnesium. 


107 


Kxp.  94,  with  a  pill-box  of  fine  shot.  Then  remove  the  cork  and 
place  in  the  tube  some  cleaned  fragments  of  magnesium  ribbon 
(rather  less  than  half  a  gram).  Replace  the  cork  and  carefully 
weigh  the  magnesium  by  adding  weights  to  the  "  tare." 

Support  the  tube  on  a  stand  as  in  the  figure,  and  pass  a  slow 
stream  of  oxygen  through  it ;  the  gas  being  previously  collected  in 
a  gas  holder,  or  being  used  from  a  cylinder  of  compressed  oxygen. 
In  order  to  see  the  rate  at  which  the  gas  is  passing,  it  should  be 


FIG.  69. 

made  to  bubble  through  water  in  the  manner  shown  at  \V  in  the 
figure.  Now  heat  the  magnesium,  applying  the  heat  gradually  at 
first,  and  then  holding  the  flame  so  as  to  ignite  the  metal  nearest 
to  the  cork.  The  metal  burns  with  great  brightness,  and  if  the  gas 
is  not  passing  in  too  fast,  the  plug  of  wool  will  prevent  any  of  the 
white  oxide  from  escaping.  When  all  the  metal  is  burnt,  and  the 
tube  has  cooled,  again  weigh  it,  and  calculate  the  proportions  in 
which  the  two  elements  have  combined,  as  in  the  former  experi- 
ment. If  the  result  is  the  same,  then  we  have  not  only  proved  that 
combustion  in  air  is  the  same  process  as  combustion  in  pure 
oxygen,  but  we  have  strengthened  our  suspicion  that  chemical 
combination  takes  place  only  between  fixed  proportions  of  the 
elements. 

Instead  of  repeating  these  experiments  a  great  many 
times,  in  order  to  prove  beyond  all  doubt  that  the  elements 
magnesium  and  oxygen  always  combine  in  the  same  fixed 


io8 


Simple  Quantitative  Manipulations. 


proportions,  it  will  be  better  to  experiment  with  other  sub- 
stances, and  see  if  they  also  show  indications  of  a  similar 
behaviour. 

Experiment  96. — The  combination  of  sulphur  with  oxygen. 
Take  a  piece  of  combustion  tube  as  in  Exp.  95,  and  fit  a  cork  and 
tube  into  each  end.  Fit  up  a  U  tube  with  corks  and  bent  tubes  as 
shown  in  Fig.  70,  attaching  a  thin  wire  to  it  by  which  it  can  be 
suspended.  Place  in  this  tube  a  small  quantity  of  a  solution  of 
sodium  hydroxide,  so  as  just  to  fill  the  bend. 

Now  counterpoise  the  combustion  tube  with  shot,  then  place  a 
small  fragment  of  sulphur  (from  ^  to  f  gram)  in  the  tube  and 
carefully  weigh  it.  Note  the  weight,  and  remove  the  "  tare  "  from 


FIG.  70. 

the  balance,  placing  it  upon  a  clean  sheet  of  paper  upon  which  can 
be  notified  what  the  tare  belongs  to. 

Next  suspend  the  U  tube  from  the  hook  at  the  end  of  the  arm 
of  the  balance,  and  counterpoise  this  with  another  tare,  carefully 
preserving  the  latter  along  with  the  former  one. 

The  U  tube  is  now  to  be  attached  to  the  combustion  tube,  as 
shown  in  the  figure,  and  a  slow  stream  of  oxygen  passed  through 
the  apparatus  using  the  bottle  W,  as  in  Exp.  95.  By  comparing 
the  rate  at  which  the  gas  bubbles  through  the  wash-bottle  and  the 
U  tube,  it  will  be  easy  to  see  whether  the  connections  are  all  tight. 
Now  gently  heat  the  sulphur.  It,  of  course,  first  melts,  and  hence 
it  is  necessary  that  the  tube  shall  be  quite  horizontal.  Presently 
it  takes  fire  and  burns,  producing  sulphur  dioxide,  the  same 


The  Combustion  of  Sulphur.  109 

compound  as  was  formed  in  Exp.  6|.  This  oxide,  as  we  then  saw,  is 
a  gas,  and  in  order  to  catch  it,  we  are  now  making  it  bubble 
through  the  solution  of  sodium  hydroxide  in  the  U  tube.  This 
readily  absorbs  sulphur  dioxide.  Notice  that  as  soon  as  the  sulphur 
begins  to  burn,  the  bubbling  in  the  U  tube  gets  slower  and  (pro- 
vided the  stream  of  oxygen  is  not  too  fast)  soon  stops  altogether, 
although  it  continues  bubbling  through  the  bottle  at  the  same 
rate  as  before.  By  the  time  the  sulphur  has  burnt  away,  however, 
it  will  be  seen  that  gas  once  more  passes  out  through  the  U  tube. 
Continue  the  stream  of  oxygen  for  a  minute  or  two  after  the  sulphur 
is  extinguished,  in  order  to  be  sure  that  all  the  gaseous  oxide  has 
been  driven  into  the  U  tube,  and  then  stop  the  experiment. 

Now  disconnect  the  U  tube,  hang  it  on  the  balance  ana  place 
its  tare  in  the  other  scale.  Observe  that  the  apparatus  has  gained 
weight :  find  how  much  by  adding  weights.  Note  the  weight. 
Next  place  the  combustion  tube  on  the  balance,  with  its  own  tare, 
and  see  if  it  is  still  counterpoised.  If  it  is,  then  the  whole  of  the 
sulphur  has  been  burnt ;  but  if  not,  add  weights  to  find  the  exact 
weight  of  what  is  left  in  the  tube,  and  note  this  weight.  We  have 
now  got  the  weight  of  sulphur  that  has  been  burnt,  and  the  weight 
of  the  oxide  of  sulphur  that  has  been  produced. 

For  example — 

Suppose  the  weight  of  sulphur  put  into  tube  =  0*62  grams 
And  weight  of  residual  sulphur  left  in  tube  =  0*0 1      „ 

Then  the  actual  weight  of  sulphur  burnt  =  0*6 1  grams 

And  suppose  the  increase  of  weight  of  the  U  tube  =  1*22  grams  ; 
then,  subtracting  the  weight  of  sulphur  burnt  from  the  weight  of 
sulphur  dioxide  obtained,  we  get  T22 

o'6i 

0*61    grams  =  weight   of  oxygen 

which  has  combined  with  o'6i  grams  of  sulphur.  Therefore  the 
two  elements  oxygen  and  sulphur  have  combined  together  in  equal 
proportions,  I  part  of  oxygen  to  I  part  of  sulphur. 

Just  as  in  the  case  of  magnesium  and  oxygen,  if  we  were 
to  repeat  this  experiment  a  number  of  times,  and  each  time 
get  the  same  result,  we  should  be  forced  to  the  conclusion  that 
sulphur  and  oxygen  do  not  unite  together  in  any  proportions, 


no  Simple  Quantitative  Manipulations. 

but  that  there  is  some  fixed  relation  between  the  weights  of 
each  which  enter  into  combination. 

If  we  repeat  this  experiment,  using  air  (contained  in  a  gas- 
holder) instead  of  oxygen,  we  shall  be  able  to  confirm  the 
former  conclusion  we  have  come  to  (p.  107),  viz.  that  the 
same  products  of  combustion  are  formed  whether  air  or  pure 
oxygen  be  employed,  and  we  shall,  at  the  same  time,  get  a 
second  result  to  serve  as  a  check  upon  the  first. 

In  carrying  out  this  experiment,  it  will  be  instructive  to 
notice  that,  during  the  burning  of  the  sulphur,  there  is  scarcely 
any  perceptible  change  in  the  rate  at  which  bubbles  pass 
through  the  liquid  in  the  U  tube.  This  is  because  in  the  air 
the  oxygen  is  mixed  with  a  large  excess  (four  times  as  much) 
of  nitrogen,  which  takes  no  part  in  the  experiment,  but  simply 
passes  through  the  apparatus  unaffected.  Both  in  the  case  of 
magnesium  and  oxygen,  and  sulphur  and  oxygen,  we  have 
been  dealing  with  instances  of  the  simple  union  of  two 
elements ;  let  us  take  one  more  illustration  of  the  synthetical 
formation  of  a  compound,  but  this  time  produced  by  indirect 
processes. 

We  have  learnt  from  Exp.  80  (p.  74)  that  hydrogen  and 
oxygen  can  be  made  to  combine  indirectly  by  first  uniting  the 
oxygen  to  copper  (forming  copper  oxide),  and  then  reducing 
this  compound  by  means  of  hydrogen,  which  deprives  the 
copper  oxide  of  oxygen,  leaving  metallic  copper.  We  shall 
now  repeat  this  experiment  quantitatively. 

Experiment  97. — The  combination  of  hydrogen  with 
oxygen.  Prepare  a  piece  of  combustion  tube,  as  in  Exp.  95,  with  a 
cork  and  straight  tube  fitted  into  one  end,  and  place  in  the  tube  a  roll 
of  clean  copper  gauze  (one  of  the  bright  rolls  after  Exp.  80  will  answer 
best).  Heat  the  tube  containing  the  copper,  with  a  Bunsen  flame, 
and,  while  it  is  hot,  pass  a  stream  of  oxygen  through  it,  so  as  to 
convert  a  portion  of  the  metal  into  its  oxide.  Then  allow  the  tube 
to  cool,  place  it  with  its  contents  upon  the  balance,  and  counter- 
poise it  with  shot.  Carefully  preserve  this  tare.  Now  attach  to 
the  other  end  of  the  tube  the  little  collecting  apparatus,  A,  Fig.  71. 
This  consists  of  a  test-tube  fitted  with  a  cork  and  two  bent  tubes. 
The  longer  one  is  drawn  out  to  a  thin  taper  and  reaches  nearly  to 
the  bottom  of  the  tube,  while  its  other  end  passes  through  a  cork 


The  Combination  of  Hydrogen  and  Oxygen.      1 1 1 

\\hich  will  properly  fit  the  end  of  the  combustion  tube.  Strong 
sulphuric  acid  is  poured  into  the  test-tube,  to  the  depth  of  about  one 
third,  care  being  taken  that  the  acid  does  not  wet  the  neck  of  the 
tube  where  the  cork  fits,  and  the  apparatus  is  then  suspended  by  a 
wire  upon  the  balance,  and  counterpoised  with  shot ;  the  tare  being 
carefully  preserved.  As  soon  as  it  has  been  counterpoised,  a  little 
cap  should  be  slipped  upon  the  end  E  of  the  short  tube,  in  order  to 
prevent  moisture  from  the  air  getting  in  when  it  is  not  being  used. 
(The  cap  is  made  by  means  of  a  short  piece  of  indiarubber  tube, 
with  a  piece  of  glass  rod  pushed  into  one  end,  as  is  shown  at  C, 
Fig.  71.)  When  this  is  attached  to  the  combustion  tube,  the  little 
cap  is  removed,  and  a  slow  stream  of  hydrogen  is  passed  through 


FIG.  71. 

the  apparatus  ;  the  tube  at  the  same  time  being  strongly  heated. 
As  what  we  are  ultimately  to  collect  and  weigh  is  water,  it  is  very 
necessary  that  the  hydrogen  we  use  should  not  bring  any  moisture 
into  the  apparatus,  or  this  would  spoil  the  result ;  the  hydrogen, 
therefore,  before  it  enters  the  tube  containing  the  copper  oxide  must 
be  made  to  pass  first  through  an  arrangement  (w,  Fig.  71)  similar 
to  the  one  we  are  using  to  absorb  the  water  formed  in  the  experi- 
ment, where  it  bubbles  through  strong  sulphuric  acid,  from  a  fine 
drawn  out  tube.  When  the  copper  in  the  tube  appears  bright 
again,  remove  the  lamp,  and  allow  the  apparatus  to  cool,  while 
the  hydrogen  still  slowly  passes  through.  Then  stop  the  gas, 
disconnect  the  collecting  apparatus,  and  place  it  en  the  balance 
with  its  own  tare.  It  is  of  course  heavier  than  before,  on  account 


I  12 


Simple  Quantitative  Manipulations. 


of  the  water  it  has  absorbed.  Carefully  weigh  the  increase,  and 
note  down  the  weight.  Now  place  the  combustion  tube  on  the 
balance,  with  its  tare ;  notice  that  it  has  lost  in  weight.  Find  how 
much  loss  by  putting  weights  in  the  scale  containing  the  tube. 
Note  this  weight.  The^tfzVz  in  weight  of  the  collecting  tube,  is  the 
weight  of  water  formed  in  the  experiment.  The  loss  of  weight 
suffered  by  the  combustion  tube,  is  the  oxygen  which  has  been 
taken  away  from  the  copper  oxide. 

Suppose   the  weight   of  water   formed    was    r2$    grams,    and 
the  weight  of  oxygen  from  the  copper  oxide    rii    grams;  then, 
subtracting  the  weight  of  oxygen  from  the  weight  of  water,  we  get 
1-25 
i-ii 

0-14  grams  =  weight  of  hydrogen  which  has  combined  with  rir 
grams  of  oxygen. 

Then,  as  0*14  :  rn  : :  i  :  x 
^  ri^  i 
0-14 

Therefore  hydrogen  and  oxygen  have  combined  together  in  the 
proportion  of  I  part  of  hydrogen  to  7-9  (or  nearly  8)  parts  of 
oxygen. 

This  experiment  has  been  made  a  number  of  times  by 
different  chemists,  with  much  more  care  and  a  great  many 
more  precautions  against  possible  errors.  Fig.  72  shows  the 


FIG.  72. 


famous  apparatus  employed  by  Dumas  (1843),  who  made  many 
very  exact  determinations  of  the  proportion  by  weight  in  which 
oxygen  and  hydrogen  combine.  In  essence  it  is  the  same  as 
that  which  has  been  used  in  the  last  experiment.  The  weighed 


The  Composition  of  Potassium  Chlorate.          1  1  3 

copper  oxide  was  heated  in  the  bulb  A-  Hydrogen  was 
generated  in  the  bottle  H,  and  the  water  formed  was  collected 
in  the  bulb  g.  The  U  tubes  i  to  8  were  to  remove  all 
impurities  from  the  hydrogen,  and  render  it  absolutely  dry  ; 
tubes  9,  10,  n,  were  to  absorb  any  water  vapour  which  did 
not  condense  in  the  bulb  B,  and  were  therefore  weighed  before 
and  after  the  experiment,  and  any  increase  in  weight  was 
added  to  the  weight  of  the  water  in  B.  Tube  12  was  merely 
a  guard  tube,  to  prevent  any  moisture  from  the  air  getting  into 
tube  ii. 

If  it  is  true  that  elements  can  combine  together  only  in 
some  fixed  proportions,  it  must  follow  that  any  particular 
compound  will  always  have  the  same  composition.  Let  us 
now  test  the  matter  from  this  point,  and  ascertain  whether 
the  same  compound  always  does  contain  its  constituent 
elements  in  the  same  proportions. 

Experiment  98.  —  The  decomposition  of  potassium  chlorate. 
Exp.  59  has  taught  us  that  this  compound  contains  oxygen  as  one 
of  its  constituent  elements  ;  and  that  when  it  is  heated  it  gives  up 
this  oxygen,  and  that  a  residue  of  potassium  chloride  is  left. 

Counterpoise  a  porcelain  crucible  with  its  lid,  and  weigh  into  it 
a  small  quantity  (say  2  or  3  grams)  of  potassium  chlorate.  Heat 
the  crucible,  as  in  Exp.  92,  keeping  the  lid  on,  because  the  salt 
decrepitates  (that  is,  "  crackles  ")  when  first  heated,  and  some  might 
be  lost  if  the  lid  were  off.  When  all  effervescence  is  at  an  end,  the 
crucible  is  allowed  to  cool,  and  is  then  weighed.  The  loss  of 
weight  is  the  oxygen  which  has  been  given  up. 

Suppose  the  weight  of  potassium  chlorate  used  to  have  been 
2'34  grams,  and  the  weight  of  the  residue  after  heating  to  have 
been  1*42  grams,  then  the  difference,  namely,  0*92  grams  =  weight 
of  oxygen  evolved,  therefore  2*34  grams  of  potassium  chlorate 
were  made  up  of  — 

Potassium  chloride  =  i  '42  grams 
oxygen  =  0-92        „ 


or,  calculating  as  in  the  former  cases,  we  get  the  proportion- 

oxygen  i  part  to  potassium  chloride  1*54  parts. 
If  on  repeating  this   experiment,   using    various    samples    of 


114  Simple  Quantitative  Manipulations. 

potassium  chlorate,  the  same  proportion  is  always  found  to  exist, 
we  should  infer  that  this  compound  always  has  the  same  com- 
position. 

How  to  calculate  the  Percentage  Composition. — 

Sometimes  chemists  prefer  to  express  the  results  of  such 
experiments  as  the  foregoing,  in  parts  per  hundred ;  that  is  to 
say,  to  state  how  much  of  the  different  ingredients  there  are 
in  100  parts  of  the  compound.  This  percentage  composition  of 
a  compound  is  calculated  from  the  experimental  numbers  in 
the  following  way,  taking  the  last  experiment  as  an  example. 

If  2*34  grams  of  potassium  chlorate  contain  1*42  grams 
of  potassium  chloride,  what  will  100  parts  contain? 

As  2-34  :  ICQ  : :  1-42  :  x 
g=r42XioQ=,6 
2'34 

Again,  if  2-34  grams  of  potassium  chlorate  contain  0*92 
grams  of  oxygen,  what  will  100  parts  contain? 

As  2*34  :  100  : :  0*92  :  x 
x  =  °:9^X*oo  = 
2-34 

Therefore  the  percentage  composition  of  potassium  chlorate, 
as  calculated  from  the  results  obtained  in  Exp.  98, x  is — 

Potassium  chloride  =  607 
oxygen  =39-3 

lOO'O 

Experiment  99. — The  decomposition  of  mercuric  oxide. 
Counterpoise  a  hard  glass  test-tube,  and  weigh  into  it  a  small 
quantity  of  mercuric  oxide.  Note  the  weight.  Now  heat  the  red 
oxide  in  the  bottom  of  the  tube,  until  it  has  entirely  decomposed 
into  oxygen  (which  escapes),  and  mercury  which  sublimes  upon 
the  upper  part  of  the  tube  ;  that  is,  until  nothing  is  left  in  the 

1  The  student  should  calculate  in  a  similar  manner,  from  his  results 
obtained  in  the  preceding  experiments,  the  percentage  composition  of 
magnesium  oxide,  of  sulphur  dioxide,  and  of  water. 


The  Composition  of  Mercuric  Oxide.  1 1  5 

bottom  of  the  tube.  When  cold,  weigh  again ;  this  gives  the 
weight  of  the  mercury,  and  the  difference  between  this  and 
the  former  weight  is  the  oxygen  which  has  been  expelled.  For 
example — 

Suppose  the  weight  of  mercuric  oxide  to  have  been  5 '40  grams, 
and  the  weight  of  mercury  left  in  the  tube  5*00  grams, 
then  5*40 
5'QQ 

0*40  grams  =  weight  of  oxygen  which  was  expelled  from  5*40 
grams  of  the  oxide. 

Calculated  as  in  the  former  cases,  we  find  that  the  proportion 
of  oxygen  to  mercury  is  as  i  to  I2'5. 

To  express  the  percentage  composition  of  the  compound,  we 
have — 

5*40  :  loo  ::  5-00  :  x  =  92-59  mercury 
and  5-40  :  100  : :  0-40  :  x  =    7 '4 1  oxygen 


lOO'OO 


In  both  of  the  last  experiments  we  have  ascertained  the 
amount  of  oxygen  by  difference,  assuming  the  loss  suffered  on 
heating  the  compounds  to  be  entirely  oxygen.  It  will  be 
obvious,  however,  that  we  must  be  quite  sure  that  nothing 
else  but  oxygen  is  expelled  by  heating.  Suppose,  for  instance, 
that  either  the  potassium  chlorate  or  the  mercuric  oxide  con- 
tained some  moisture ;  this  would  be  driven  off  along  with 
the  oxygen,  and  the  loss  of  weight  due  to  this  would  be  counted 
as  oxygen.  We  can  check  the  former  result  by  modifying 
the  experiment  so  as  to  collect  and  measure  the  oxygen,  and 
then  by  the  methods  of  calculation  explained  on  page  101, 
we  can  find  from  the  volume  of  gas  collected  what  its  weight 
is,  without  actually  weighing  it. 

Experiment  100. — Fit  a  short,  narrow,  hard  glass  test-tube  (or 
piece  of  combustion  tube  closed  at  one  end)  with  a  cork  and 
narrow  delivery  tube.  Counterpoise  the  empty  tube,  and  weigh 
into  it  a  small  quantity  of  mercuric  oxide.  Attach  the  delivery 
tube,  and  proceed  to  heat  the  compound,  collecting  the  oxygen 
over  water  in  the  usual  way  in  a  cylinder,  capable  of  holding  about 
a  litre.  As  soon  as  the  whole  of  the  oxide  is  decomposed,  dis- 
connect the  delivery  tube,  and,  when  cold,  again  weigh  the  tube,  to 


n6  Simple  Quantitative  Manipulations. 

find  the  weight  of  the  mercury.  Now  mark  the  volume  of  gas  in 
the  cylinder  x  by  sticking  a  strip  of  a  gummed  label  upon  the  glass. 
Take  the  .temperature  of  the  water  in  the  trough  by  dipping  a 
thermometer  into  it ;  this  may  be  taken  as  the  temperature  of  the 
gas.  To  find  the  exact  pressure  under  which  the  gas  is,  we  must 
measure  the  height  of  the  water  in  the  cylinder  above  the  level  of 
the  liquid  in  the  trough,  and  divide  this  by  13*6  (because  mercury 
is  13*6  times  as  heavy  as  water)  ;  and  then  subtract  the  result  from 
the  height  of  the  barometer  at  the  time.  '(If,  however,  the  cylinder 
used  is  not  a  long  one,  and  if  the  atmospheric  pressure  is  not  very 
far  from  the  normal,  viz.  760  mm.,  the  final  result  will  not  be  much 
influenced  by  this  correction,  and  it  can  be  omitted.) 

Now  remove  the  cylinder  from  the  trough,  and  empty  it.  Then 
pour  water  into  it  from  a  graduated  cc.  measure,1  until  the  liquid 
reaches  to  the  edge  of  the  paper  mark.  The  number  of  cc.s 
required  to  do  this  will  give  the  volume  of  oxygen,  measured  at 
the  observed  temperature  and  pressure. 

In  the  following  example,  correction  has  been  made  for  pressure 
as  well  as  temperature. 

Weight  of  mercuric  oxide  taken  =  8 '90  grams 

„        „      mercury  remaining  =  8^24     „ 
Number  of  cc.s  of  water  required  to  fill  the  cylinder  \ 

up  to  the  label !  =  485  cc. 

That  is,  the  volume  of  oxygen  as  collected       .        . ) 

Temperature  =  16°  C. 

Pressure — Height  of  water  in  cylinder,  above  level  in  trough 

=  90  millimetres 

Corresponds  to  •&.#  =  6*6  millimetres  of  mercury 
Barometer  reading  =  772*6 
.'.  Actual  pressure  =  772*6  —  6*6  =  766  mm. 

(1)  485  cc.  of  oxygen,  at  16°  and  766  mm.,  what  volume  at 
N.T.P.  ? 

485x273x766    =    6      cc 
(273  +  16)  x  760 

(2)  What  is  the  weight  of  4617  cc.  of  oxygen  £ 

loco  cc.  (i  litre)  of  oxygen  weighs  0-0896  x  16  =  i'4336  grams 
(see  p.  101). 

.*.  looo  :  4617  : :  r4336  :  x  =  o'66  grams. 

1  If  a  graduated  cc.  cylinder  is  used  in  which  to  collect  the  gas,  the 
volume  can  be  read  oft'  at  once,  and  the  opfiration  of  measuring  the 
capacity  of  the  cylinder  will  of  course  be'unnecessary. 


The  Law  of  Constant  Composition.  117 

Therefore  '66  grams  of  oxygen  were  obtained  from  8*90 
grams  of  mercuric  oxide.  ' 

Or  oxygen  and  mercury  are  present  in  the  proportion  of  I  :  12*5. 

First  Law  of  Chemical  Combination. — By  making  a 
large  number  of  experiments  like  these,  we  are  driven  to  the 
general  conclusion  that  elements  do  not  combine  in  any  pro- 
portion, but  that  they  always  unite  in  some  definite  and  fixed 
quantities.  We  have  in  this  way  discovered  a  general  truth, 
or  a  Law,  as  it  is  termed — the  Law  of  Constant  Proportion, 
which  may  be  stated  in  the  following  words,  The  same  com- 
pound always  contains  the  same  elements,  combined  together  in 
the  same  proportion  by  weight. 

Hydrogen  and  oxygen,  we  have  seen,  unite  in  the  pro- 
portion by  weight  of  i  part  of  hydrogen  to  8  parts  of  oxygen. 
If,  therefore,  these  two  gases  be  mixed  together  in  any  other 
proportion  before  combination,  there  will  remain  over  after 
their  union  the  excess  of  either  the  one  or  the  other  as  the 
case  may  be.  Thus,  if  equal  weights  were  present  in  the 
mixture  before  chemical  union,  when  combination  took  place 
the  oxygen  would  still  only  combine  with  -|th  of  its  weight 
of  the  hydrogen,  leaving  the  remaining  J-  ths  of  the  hydrogen 
unaffected.  Similarly,  if  any  excess  of  oxygen  be  present 
beyond  the  proportion  of  i  part  of  hydrogen  to  8  parts  of 
oxygen,  this  excess  remains  over  unaltered  after  combination 
of  the  two  elements. 

Oxygen  and  magnesium  again  combine  in  the  proportion 
of  i  to  1-5.  Any  excess  of  either  element  beyond  these  pro- 
portions that  may  be  present  before  combination,  is  left  behind 
unaltered  when  the  elements  unite. 

Second  Law  of  Chemical  Combination. — In  the 
course  of  experimenting  with  a  view  to  establish  the  law  of 
constant  proportion,  chemists  meet  with  a  number  of  cases 
which  at  first  appear  like  exceptions,  for  it  is  found  that  very 
often  two  elements  do  combine  together  in  more  than  one 
proportion.  Whenever  this  happens,  however,  the  compounds 
that  are  produced  are  totally  different  from  each  other  j  and 
moreover  the  proportion  in  which  the  elements  are  present  in 
each,  is  found  to  be  always  the  same  for  the  same  compound  ; 


Ii8  Simple  Quantitative  Manipulations. 

that  is  to  say,  each  compound  has  its  own  constant  composi- 
tion. For  instance,  we  have  found  that  when  sulphur  burns 
in  air  or  in  oxygen,  an  oxide  of  sulphur  is  formed  which 
contains  the  two  elements  in  the  proportion  S  :  O  =  i  :  i,  and 
we  have  learnt  that  the  composition  of  this  compound  is  always 
the  same.  But  under  certain  conditions,  sulphur  and  oxygen 
can  be  made  to  unite  in  a  different  proportion,  namely, 
S  :  O  =  i  :  i  '5.  The  compound  that  is  obtained  under  these 
circumstances,  however,  is  a  totally  different  substance  from 
the  other ;  its  composition  is  always  absolutely  constant,  never 
departing  from  the  proportion  of  i  part  sulphur  to  i-|  parts  of 
oxygen.  Similarly  the  element  carbon  combines  with  oxygen 
in  two  proportions,  giving  rise  to  two  distinct  compounds,  one 
of  them  contains  the  elements  in  the  proportion  of  carbon 
i  part,  and  oxygen  i  "33  parts,  while  the  other  is  composed  of 
carbon  i  part,  and  oxygen  2-66  parts;  that  is,  it  contains  twice 
as  large  a  proportion  of  oxygen.  The  composition  of  each 
compound  never  varies,  and  therefore  they  are  not  exceptions 
to  the  law  of  constant  proportions. 

Another  striking  example  of  the  same  elements  combining 
in  more  than  one  proportion  is  the  case  of  oxygen  and 
nitrogen.  These  two  elements  unite  in  no  less  than  five 
different  proportions,  giving  five  different  compounds,  each 
one  quite  different  from  the  rest,  and  each  having  a  perfectly 
constant  composition.  Their  names  and  composition  are  as 
follows : — 


Nitrous  oxide 

Nitric  oxide    


Nitrogen  trioxide  ... 
Nitrogen  peroxide ... 
Nitrogen  pentoxide 


part  of  nitrogen  to  0*57  parts  of  oxygen. 


1-14 

171 
2-28 
2-85 


That  is,  the  second,  third,  fourth,  and  fifth  of  the  compounds 
contain  respectively  twice,  three  times,  four  times,  and  five 
times  as  much  oxygen  in  their  composition  as  the  first ;  or  the 
relative  proportions  of  oxygen  which  are  combined  with  a 
fixed  quantity  of  nitrogen  are  as  i  :  2  :  3  :  4  :  5 

In  all  cases  like   these,  where  the  same  elements  com- 
bine in  more  than  one  proportion,  forming  more  than  one 


Law  of  Multiple  Proportions.  119 

compound,  we  find  the  same  simple  ratio  existing  between  the 
different  weights  of  the  element  which  can  combine  with  a 
constant  weight  of  the  other.  This  is  another  general  truth, 
or  law,  known  as  the  Law  of  Multiple  Proportions,  which  may 
be  thus  stated — 

When  the  same  two  elements  unite  to  form  more  than  one 
compound,  the  different  weights  of  the  one,  that  combine  with  a 
constant  weight  of  the  other \  bear  a  simple  ratio  to  one  another. 


CHAPTER   XII. 

HYDROCHLORIC   ACID. 

IN  beginning  the  study  of  this  substance,  the  first  point  to 
settle  is  whether  it  is  an  element  or  a  compound,  and  if  the 
latter,  what  elements  it  is  composed  of.  We  have  already 
learnt  (p.  47),  that  when  this  substance  is  acted  upon  by  zinc, 
hydrogen  is  evolved.  It  must,  therefore,  contain  hydrogen, 
and  therefore  it  is  a  compound ;  but  what  is  the  hydrogen 
combined  with  ?  What  else  does  the  compound  contain 
besides  hydrogen  ? 

The  only  other  compound  we  have  examined  so  far,  is 
water ;  let  us,  therefore,  try  to  decompose  hydrochloric  acid 
as  water  was  decomposed,  by  an  electric  current. 

Experiment  101. — Place  some  hydrochloric  acid  in  the  apparatus 
(Fig.  56,  p.  76)  and  connect  the  wires  to  a  battery.  Collect  all  the 
gas  in  a  short  stout  tube,  as  in  Exp.  81.  When  the  tube  is  full, 
apply  a  light  to  it.  Notice  that  it  explodes  very  much  as  the  gas 
obtained  from  water  did.  Can  it  be  oxygen  and  hydrogen  ? 

Experiment  102. — Collect  the  gas  from  each  wire  in  separate 
tubes,  as  in  Exp.  82.  Notice  that  the  two  gases  are  not  evolved  in 
the  proportions  they  were  when  water  was  decomposed.  In  that 
case  there  was  twice  as  much  hydrogen  as  oxygen,  here  the  two 
gases  collect  in  nearly  equal  volumes.1  Test  the  gas  which  collects 
from  the  wire  which  is  connected  with  the  zinc  (or  negative)  plate 
of  the  battery,  by  removing  the  test-tube  and  applying  a  light  to 
the  gas.  Note  that  it  burns  in  the  familiar  manner  of  hydrogen. 
Notice  that  the  gas  in  the  other  tube  is  not  quite  colourless,  but  is 
faintly  yellow.  Cautiously  bring  a  lighted  taper  into  the  tube  and 

1  If  the  current  is  passed  through  the  acid  for  some  little  time  before 
the  gases  are  collected,  they  will  collect  in  equal  volumes. 


Hydrochloric  Acid. 


121 


observe  the  behaviour  of  the  flame.  It  does  not  burn  brilliantly, 
therefore  the  gas  is  certainly  not  oxygen.  Neither  is  the  flame 
altogether  extinguished,  but  it  burns  with  a  strangely  lurid  and 
smoky  manner. 

Once  more  connect  the  battery  for  a  few  moments,  and  gently 
sniff  the  gas  which  is  evolved,  observe  that  it  has  a  most  irritating, 
cough-producing  effect  upon  the  air  passages  of  the  nose  and  upon 
the  throat ;  quite  different  from  the  smell  of  the  hydrochloric  acid 
itself.  This  is  a  new  gas,  for  neither  oxygen  nor  hydrogen  have 
any  smell,  or  colour,  or  behave  towards  a  taper-flame  as  this  gas 
does.  This  yellowish,  choking  gas  is  called  chlorine ;  the  name 
being  from  the  Greek  word  for  yellowish  green. 

In  the  case  of  water,  we  found,  when  the  two  gases 
obtained  by  electrolysis  were  mixed  together  in  a  strong 
glass  vessel  and  exploded,  that  a  small  quantity  of 
liquid  water  was  formed  and  that  there  was  no  gas  left 
in  the  vessel  (Cavendish's  experiment).  Let  us  repeat 
this  experiment  with  the  two  gases  given  off  when 
hydrochloric  acid  is  submitted  to  electrolysis. 

Experiment  103. — Fit  a  small  bottle  with  a  cork  carrying 
a  glass  tube  bent  at  a  right  angle.  Push  through  the  cork 
two  brass  (or,  better,  platinum)  wires,  the  ends  of  which 
are  twisted  round  two  rods  of  carbon  (as 
in  Fig.  73)  about  the  thickness  of  a  slate 
pencil  (such  carbon  rods  as  are  used  for 
electric  lamps).  Fill  the  bottle  with  hydro- 
chloric acid  nearly  up  to  the  wires,  and 
connect  the  wires  to  a  battery.  Allow  the 
current  to  pass  for  about  ten  minutes. 
Then  connect  to  the  exit-tube  a  long 
stout  glass  tube  (Fig.  74),  having  a  stop- 
cock at  each  end,  and  provided  with  two 
platinum  wires  sealed  into  the  glass. 
Allow  the  gases  to  stream  through  this 
tube  for  ten  to  fifteen  minutes  (depending 
on  the  rate  at  which  the  hydrochloric  FlG<  74' 

acid  is  being  decomposed  ;  that  is,  on  the  strength  of  battery 
being  used),  then  close  both  cocks  and  stop  the  current.  Now, 
by  means  of  a  Ruhmkorf  coil,  send  an  electric  spark  into  the 
gases  in  the  tube  by  the  wires.  The  gases  instantly  explode, 
but  notice  that  there  is  no  noise,  only  a  slight  click  accompanied 


122  Hydrochloric  Acid. 

by  a  flash  of  light  in  the  tube.  Carefully  examine  the  tube  to  see 
if  there  is  any  indication  of  a  liquid  moistening  the  inside.  There 
is  none.  If  the  hydrochloric  acid  which  has  been  formed  by  the  re- 
union of  the  gases  has  not  condensed  to  a  liquid,  where  is  it  ?  Dip 
one  end  of  the  tube  beneath  some  mercury  in  a  dish,  and  open  the 
stop-cock  at  that  end.  Notice  that  the  mercury  is  not  sucked  in,  and 
also  that  no  gas  bubbles  out.  This  shows  that  the  tube  is  exactly 
full  of  gas,  just  as  it  was  at  first ;  therefore,  if  hydrochloric  acid  has 
been  produced  it  must  be  a.  gas  and  not  a  liquid,  and  also  it  must 
"occupy  exactly  as  much  space  or  volume  as  the  mixture  of  gases 
that  were  put  into  the  tube.  Now  open  one  end  of  the  tube  beneath 
water.  Notice  that  the  water  is  sucked  up  into  the  tube  exactly 
as  though  the  tube  were  vacuous  !— but  it  was  not  vacuous,  because, 
had  it  been  so,  the  mercury  would  have  entered  before.  Test  a 
little  of  the  water  which  has  been  drawn  up  into  the  tube,  with  a 
piece  of  litmus  paper  ;  notice  that  the  water  is  acid.  The  hydro- 
chloric acid  in  the  gaseous  state  which  was  formed  by  the  union  of 
the  gases  in  the  tube  has  dissolved  in  the  water. 

We  see  from  this,  that  hydrochloric  acid  is  a  gas  at  ordinary 
temperatures.  What,  then,  is  the  liquid  we  have  used,  and 
which  is  generally  called  hydrochloric  acid.  This  liquid  is  simply 
a  strong  solution  of  the  gas  in  water.  In  Exp.  103  we  have 
obtained  a  small  quantity  of  the  gas  by  direct  synthesis.  Let 
us  now  prepare  a  larger  quantity  of  it  by  an  indirect  method, 
and  study  its  properties.  To  produce  it,  we  shall  use  sul- 
phuric acid  and  sodium  chloride.  We  know  that  the  first 
of  these  contains  hydrogen;  the  sodium  chloride  (common 
salt)  contains  chlorine  combined  with  sodium.  When  these 
two  compounds  are  brought  together  a  chemical  change  takes 
place.  The  sodium  of  the  sodium  chloride  displaces  the 
hydrogen  of  the  sulphuric  acid ;  and  this  hydrogen  then 
unites  with  the  chlorine  which  the  sodium  has  left. 

Experiment  104. — Measure  55  cc.  of  strong  sulphuric  acid  into 
a  flask,  and  slowly  pour  upon  it  40  cc.  of  water,  constantly  shaking 
the  mixture.  The  liquids  become  extremely  hot,  and  must  be 
allowed  to  cool.  Fit  the  flask  with  a  cork  carrying  a  straight  tube 
which  just  dips  into  the  liquid,  and  a  short  bent  exit  tube.  Put  a 
handful  of  common  salt  into  the  cold  acid,  attach  a  delivery  tube 
and  prepare  to  collect  the  gas  in  the  usual  way.  Little  action  takes 
place  in  the  flask  in  the  cold,  but  on  gently  warming  it  a  plentiful 


Solubility  of  Hydrochloric  Acid.  123 

evolution  of  gas  takes  place.  At  first  the  air  within  the  apparatus 
is  expelled  and  bubbles  through  the  water  in  the  trough  ;  but  notice 
that  presently  the  bubbles  get  smaller  and  smaller,  until  scarcely  a 
trace  of  gas  escapes  through  the  water.  Why  is  this  ?  Evidently 
gaseous  hydrochloric  acid  cannot  be  collected  over  water.  Now 
disconnect  the  delivery  tube,  and  collect  a  cylinder  of  the  gas  by 
downward  displacement  (p.  27).  Notice  that  the  gas  is  quite 
colourless,  but  that  when  it  escapes  into  the  air  it  fumes  strongly. 
Gently  fan  a  whiff  of  the  gas  towards  the  face,  so  as  to  smell  it  ; 
and  note  that,  although  very  pungent,  it  is  quite  different  from 
chlorine.  When  the  cylinder  is  full,  cover  it  with  a  stout  glass 
plate,  and  quickly  invert  it  mouth  downward  in  a  trough  of  water. 
Observe  that  the  water  quickly  rushes  up  into  the  cylinder,  as 
though  the  latter  were  vacuous,  showing  how  extremely  sohible  this 
gas  is,  and  how  utterly  impossible  it  would  be  to  collect  it  over 
water.  Test  the  water  in  the  cylinder  with  litmus,  see  that  it  is 
acid.  Although  this  gas  cannot  be  collected  over  water,  it  can  be 
collected  over  mercury,  in  a  mercury  pneumatic  trough,  for  it  has 
no  action  on  this  metal. 

Experiment  105. — Collect  a  second  cylinder  of  gas.  Dip  a 
moist  litmus  paper  into  it  ;  notice  that  the  gas  is  powerfully  acid. 
Test  the  behaviour  of  the  gas  towards  a  flame,  by  bringing  a  lighted 
taper  to  the  mouth  of  the  jar.  Observe  that  the  gas  does  not  behave 
either  like  hydrogen  or  like  chlorine  ;  for  it  does  not  burn,  and  does 
not  allow  the  taper  to  continue  burning  even  with  a  smoky  flame. 

Let  us  now  try  by  chemical  processes  to  expel  first  the 
hydrogen,  and  then  the  chlorine,  from  this  gaseous  compound 
of  these  two  elements.  To  displace  the  hydrogen  we  will 
employ  the  metal  sodium. 

Experiment  106. — Place  a  fragment  of  sodium  in  a  hard  glass 
bulb-tube  (as  in  Fig.  38,  p.  44),  and  pass  a  brisk  stream  of  gaseous 
hydrochloric  acid  over  it.  Attach  a  delivery  tube  to  the  exit,  and 
collect  any  gas  over  water.  When  the  air  has  been  swept  out  of 
the  bulb  tube,  proceed  to  heat  the  sodium,  which  quickly  melts  and 
presently  takes  fire.  The  moment  it  begins  to  burn,  gas  at  once 
collects  in  the  vessel  in  the  pneumatic  trough.  Test  this  gas  with 
a  lighted  taper,  and  recognize  hydrogen  by  the  familiar  flame. 
What  has  become  of  the  other  constituent,  namely,  the  chlorine  ? 
Break  open  the  bulb,  and  with  the  little  finger  bring  a  small  quantity 
of  the  white  residue  on  to  the  tongue.  By  the  familiar  taste  recog- 
nize common  salt,  that  is,  sodium  chloride  ;  therefore  the  chlorine 


124  Hydrochloric  Acid. 

has  combined  with  the  sodium.     The  combustion  of  the  metal  was 
in  reality  its  rapid  combination  with  chlorine. 

In  order  to  liberate  the  chlorine  from  hydrochloric  acid,  we 
may  employ  certain  compounds  rich  in  oxygen,  so  that  the 
hydrogen  in  the  acid  may  unite  with  this  to  form  water,  and 
let  the  chlorine  go  free.  Substances  belonging  to  a  class  called 
peroxides  will  do  this.  We  will  use  manganese  peroxide  (also 
called  manganese  dioxide). 

Experiment  107. — Place  a  quantity  of  manganese  dioxide  in  a 
moderately  wide  glass  tube,  supported  in  a  horizontal  position  and 
fitted  at  each  end  with  a  cork  and  glass  tube.  To  one  end  connect 
a  delivery  tube,  leading  to  the  pneumatic  trough,  and  attach  the 
other  end  to  the  apparatus  generating  the  gaseous  hydrochloric 
acid.  As  the  gas  passes  through  the  tube,  the  manganese  dioxide 
may  be  slightly  warmed  by  just  brushing  a  flame  along  the  tube. 
Notice  that  a  gas  is  again  collected  through  the  water  in  the 
trough.  Observe  also  that  moisture  collects  upon  the  inside  of  the 
horizontal  tube.  This  is  the  product  of  the  union  of  the  hydrogen 
present  in  the  hydrochloric  acid  with  oxygen  out  of  the  manganese 
dioxide.  Test  the  gas  with  a  lighted  taper  and  observe  the  same 
smoky  flame  as  in  Exp.  102.  Cautiously  whiff  the  gas  towards  the 
face  and  recognize  the  same  suffocating  smell. 

By  Exps.  102  and  103  we  have  learnt  that  hydrochloric 
acid  contains  half  its  own  volume  of  each  of  its  constituent 
elements.  So  far  as  the  volume  of  the  hydrogen  is  concerned 
(but  not  that  of  the  chlorine)  this  may  be  confirmed  in  the 
following  way.  The  U-shaped  apparatus  (Fig.  75)  is  filled  with 
gaseous  hydrochloric  acid  down  to  the  second  division,  the 
mercury  being  level  in  the  two  limbs.  We  then  fill  up  the 
open  limb  with  sodium  amalgam  (this  is  a  solution  of  sodium 
in  mercury,  made  by  pressing  pieces  of  sodium  beneath  the 
surface  of  mercury  in  a  mortar  with  the  pestle.  The  sodium 
dissolves  in  the  mercury  with  great  energy,  sometimes  even 
taking  fire).  The  tube  is  then  closed  with  the  thumb,  and  tipped 
up  so  as  to  make  the  gas  pass  round  the  bend  and  bubble  up 
through  the  amalgam.  The  gas  is  passed  through  the  amalgam 
once  or  twice,  and  then  returned  to  the  closed  limb.  While  in 


Chlorides. 


125 


contact  with  the  sodium  amalgam,  the  gas  is  decomposed. 
The  chlorine  unites  with  the  sodium,  forming  sodium  chloride, 
which  appears  as  a  white  crust  on  the  mercury  and  the  sides 
of  the  tube,  while  the  hydrogen  is  set  free,  just  as  in  Exp.  106. 
When  the  mercury  in  both  limbs  is 
once  more  levelled,  by  running  away 
the  excess  by  means  of  the  side  tap, 
it  is  seen  that  the  residual  gas  occupies 
exactly  one  half  the  original  volume. 
We  can  prove  that  this  gas  is  hydro- 
gen by  again  filling  the  open  limb  with 
mercury  so  as  to  drive  the  gas  out  at 
the  stopcock  at  the  top,  and  on  apply- 
ing a  light  the  issuing  gas  will  burn  with 
the  familiar  hydrogen  flame. 

Chlorides.  —  Compounds  of  ele- 
ments with  chlorine  are  called  chlo- 
rides. Metallic  chlorides,  for  example, 
are  the  compounds  of  metals  with 
chlorine.  Many  of  them  are  produced 
by  the  direct  union  of  the  metals  with 
chlorine.  We  may,  however,  also  de- 
fine chlorides  as  compounds  in  which 
the  hydrogen  of  hydrochloric  acid  is 
replaced  by  a  metal.  Thus,  when  zinc 
is  acted  on  by  hydrochloric  acid,  the 
hydrogen  of  the  acid  is  displaced  and  FlG_  7St 

liberated   as    gas,  while    zinc   chloride 

is  at  the  same  time  produced.  Again,  if  certain  oxides  of 
metals  are  acted  on  by  hydrochloric  acid,  the  hydrogen  is 
again  displaced  and  a  chloride  of  the  metal  produced ;  but  in 
this  case  the  hydrogen  is  not  liberated  as  gas,  but  goes  to  the 
oxygen  of  the  oxide  to  produce  water.  The  same  result  is 
obtained  when  a  metallic  hydroxide  is  brought  into  contact 
with  hydrochloric  acid ;  a  chloride  of  the  metal  and  water  are 
formed.  Since  hydrochloric  acid  is  a  compound  of  hydrogen 
with  chlorine,  it  is  often  called  hydrogen  chloride. 

Tests  for  Chlorides. — Hydrogen  chloride  and  all  other 


126  Hydrochloric  Acid. 

chlorides  that  are  soluble  in  water,  are  detected  by  means  of  a 
solution  of  silver  nitrate. 

Experiment  108. — Dissolve  a  pinch  of  sodium  chloride  (common 
salt)  in  water,  and  add  a  few  drops  of  silver  nitrate  solution.  A 
white  curdy  precipitate,  consisting  of  silver  chloride,  is  at  once 
produced.  Divide  the  liquid  containing  the  precipitate  into  two 
portions,  and  add  some  ammonia  to  one  and  nitric  acid  to  the 
other.  Note  that  the  precipitate  dissolves  in  ammonia,  and  does 
not  dissolve  in  nitric  acid.  This  distinguishes  the  precipitate  from 
any  other  which  would  be  produced  by  silver  nitrate. 

Lead  nitrate  also  gives  a  similar  looking  precipitate  when 
added  to  a  solution  containing  a  chloride,  and  this  precipitate 
can  be  distinguished  from  others  in  the  following  way. 

Experiment  109. — Add  a  few  drops  of  a  solution  of  lead  nitrate 
to  a  solution  of  sodium  chloride.  Compare  the  precipitate  of  lead 
chloride  with  that  obtained  in  the  last  experiment.  Now  boil  the 
mixture,  and  notice  that  the  precipitate  disappears.  It  dissolves  in 
hot  water.  Cool  the  test-tube,  and  observe  white,  silky-looking 
crystals  begin  to  deposit.  This  is  the  lead  chloride  being  deposited 
again  in  the  form  of  crystals. 

All  chlorides,  whether  soluble  in  water  or  not,  give  off 
chlorine  when  warmed  with  a  mixture  of  manganese  dioxide 
and  sulphuric  acid,  and  can,  therefore,  be  identified  by  this 
test,  for  the  smell  and  colour  of  chlorine  are  unmistakable 
(see  Chlorine,  next  chapter). 

EPITOME. 

Hydrochloric  acid  is  a  compound  of  hydrogen  and  chlorine.  It 
is  sometimes  called  hydrogen  chloride.  It  can  be  produced  syn- 
thetically by  the  direct  union  of  the  two  constituent  elements. 

Hydrochloric  acid  is  prepared  by  acting  on  common  salt  with 
sulphuric  acid.  Under  ordinary  circumstances  it  is  a  colourless 
gas,  having  a  pungent  acid  smell.  When  it  escapes  into  the  air  it 
fumes  strongly.  It  is  extremely  soluble  in  water.  At  the  ordinary 
temperature  I  litre  of  water  will  dissolve  450  litres  of  the  gas.  This 
solution  of  the  gas  in  water  is  the  common  article  of  commerce 
known  as  hydrochloric  acid  (also  sometimes  called  muriatic  acid, 
and  spirit  of  salt}. 

Gaseous  hydrochloric  acid  is  i|-  times  as  heavy  as  air,  and  can 


Hydrochloric  Acid.  127 

therefore  be  collected  by  downward  displacement.  The  gas  does 
not  burn,  neither  will  it  support  the  combustion  of  ordinary  flames. 
It  is  strongly  acid,  but  has  no  bleaching  power. 

The  solution  of  the  gas  in  water  is  decomposed  by  a  current  of 
electricity  into  equal  volumes  of  hydrogen  and  chlorine  ;  and  when 
these  two  gases  in  equal  volumes  unite,  the  gaseous  hydrochloric 
acid  which  results  occupies  the  same  volume  as  the  mixed  gases  did 
before  their  union. 

Hydrochloric  acid  is  manufactured  on  an  enormous  scale  as  an 
article  of  commerce.  It  is  obtained  by  the  same  method  as  in  the 
laboratory,  namely,  from  salt  and  sulphuric  acid,  only  the  operation 
is  conducted  in  enormous  iron  vessels,  and  the  gas  is  made  to  pass 
at  once  into  water,  to  obtain  a  solution  of  it. 

Hydrochloric  acid  attacks  many  metals  ;  some,  such  as  zinc  and 
iron,  very  readily  ;  others,  as  copper  and  tin,  less  quickly.  In  all 
cases,  hydrogen  is  given  off,  and  a  chloride  of  the  metal  is  left. 
Hydrochloric  acid  has  no  action  upon  mercury,  or  upon  the  so-called 
"noble"  metals,  gold  and  platinum. 

Gaseous  hydrochloric  acid  is  distinguished  from  all  other  gases 
by  its  extreme  solubility  in  water,  and  by  the  formation  of.,  a  white 
curdy  precipitate  of  silver  chloride,  when  a  solution  of  silver  nitrate 
is  added  to  the  solution  of  the  gas  ;  the  precipitate  being  insoluble 
in  nitric  acid,  but  easily  soluble  in  ammonia.  All  soluble  chlorides 
are  tested  for  with  silver  nitrate  in  the  same  way. 

Reactions  for  hydrochloric  acid — 

(r)  Preparation  from  sodium  chloride  and  sulphuric  acid  (labo- 
ratory process), 

Nad  +  H2S04  =  NaHS04  +  HC1. 

(2)  Manufacturing  process— when  the  mixture  is  more  strongly 
heated,  and  all  the  hydrogen  from  the  sulphuric  acid  is  displaced 
by  sodium — 

2NaCl  +  H2SO4  =  Na2  SO4  +  2  HC1. 

(3)  Test  for  hydrochloric  acid,  with  silver  nitrate, 

HC1  +  AgNO3  =  AgCl  +  HNO3. 


CHAPTER    XIII 


CHLORINE. 


FROM  the  last  chapter  we  learn  that  chlorine  is  one  of  the 
constituent  elements  present  in  hydrochloric  acid ;  and  we  saw 
in  Exp.  107  how  it  may  be  obtained  from  this  compound. 
When  we  require  any  quantity  of  this  element,  it  is  more  con- 
venient to  start  with  a  solution  of  hydrochloric  acid  instead  of 
the  gas. 

Experiment  no. — Place  a  quantity  of  strong  hydrochloric  acid 
in  a  flask  fitted  with  a  cork  and  delivery  tube,  and  add  a  handful  of 
powdered  manganese  dioxide.  Gently  warm  the  flask,  and  collect 
the  gas  at  the  pneumatic  trough,  using  strong  brine  instead  of 
ordinary  water.  Collect  several  cylinders  full.  Notice  the  colour 
of  the  gas,  which  when  seen  in  quantity,  appears  a  greenish-yellow. 

Instead  of  employing  ready-made  hydrochloric  acid  as  the 
source  of  chlorine,  we  can  use  the  materials  from  which  hydro- 
chloric acid  is  obtained,  namely,  common  salt  and  sulphuric 
acid. 

Experiment  in. — Place  a  quantity  of  sulphuric  acid  (of  the 
same  strength  as  that  used  in  Exp.  104)  in  a  flask,  and  add  common 
salt  and  manganese  dioxide  in  about  equal  parts.  Apply  a  gentle 
heat,  and  collect  the  gas  over  strong  salt-water,  as  above. 

1  In  experimenting  with  chlorine,  the  student  must  be  careful  not  to 
inhale  the  gas.  Even  when  mixed  with  a  large  volume  of  air,  it  causes  a 
most  unpleasant  irritation  to  the  throat  and  lungs  j  and  if  a  strong  whiff 
of  the  gas  is  inhaled  it  becomes  a  dangerous  poison.  These  experiments, 
therefore,  are  best  performed  in  a  good  draught  cupboard. 


The  Bleaching  Action  of  Chlorine.  129 

Experiment  112. — Properties  of  Chlorine.  Take  one  of  the 
jars  of  gas  and  place  it  mouth  downward  in  ordinary  water.  Notice 
that  the  water  slowly  rises  in  the  jar,  and  if  allowed  to  stand  some 
time  will  gradually  creep  up  to  the  top.  This  shows  that  the  gas 
is  moderately  soluble  in  water  (contrast  this  with  the  case  of  hydro- 
chloric acid).  It  is  because  chlorine  is  thus  dissolved  by  water, 
that  in  Exp.  99  the  volumes  of  the  two  gases  were  not  quite  equal. 
There  was  rather  less  chlorine  collected  than  hydrogen.  If  when 
the  water  has  risen  a  little,  the  jar  be  closed  with  a  cork,  and  briskly 
shaken,  the  water  will  dissolve  the  gas  more  quickly,  and  it  will  be 
seen  that  the  liquid  has  much  the  same  yellowish  colour  as  the 
gas.  This  solution  is  called  chlorine  water,  and  smells  strongly 
of  chlorine.  Pour  it  into  a  stoppered  bottle,  and  stand  it  in  the 
window.  If  the  sun  happens  to  be  shining  on  the  window,  notice 
that  the  yellow  colour  rapidly  disappears,  and  at  the  same  time  the 
smell  of  the  chlorine  has  gone.  The  same  result  follows  more 
slowly  in  dull  daylight.  The  chlorine,  under  the  influence  of  light, 
decomposes  the  water  in  which  it  is  dissolved,  combining  with 
the  hydrogen  to  form  hydrochloric  acid,  while  oxygen  from  the 
water  is  set  free. 

We  have  seen  in  Exp.  103  that  chlorine  and  hydrogen 
combine  with  explosion  when  a  spark  is  passed  into  a  mixture 
of  these  gases.  The  union  of  these  elements  also  takes  place 
if  the  mixture  is  exposed  to  daylight.  It  takes  place  slowly 
and  noiselessly  in  dull  daylight,  but  instantly  with  explosion 
in  direct  sunlight  or  in  the  bright  light  of  burning  magnesium. 

Experiment  113. — Dip  a  moistened  blue  litmus  paper  into  ajar 
of  chlorine.  Notice  that  instead  of  being  reddened,  as  with  hydro- 
chloric acid,  the  paper  is  bleached.  Do  the  same  with  some  other 
colours,  such  as  carmine,  or  aniline  blue  or  violet,  by  staining  pieces 
of  paper  with  the  dyes,  and  dipping  them  into  the  gas. 

This  bleaching  power  of  chlorine  is  one  of  its  most  charac- 
teristic and  important  properties. 

Experiment  1 14.— Collect  a  stoppered  jar  of  chlorine  by  down- 
ward displacement,  having  first  placed  a  small  layer  of  strong  sul- 
phuric acid  in  the  jar.  Close  the  jar  with  a  cork,  and  spread  the 
acid  over  the  interior  as  much  as  possible  by  tipping  the  vessel 
about.  The  strong  acid  has  the  power  of  removing  any  vapour  of 

K 


1 30  Chlorine. 

water  which  is  present  in  the  gas.  Now  dry  a  strip  of  paper  which 
has  been  stained  with  carmine  or  an  aniline  colour,  by  holding  it 
for  a  few  moments  in  front  of  a  fire,  or  over  a  gas  flame.  Then, 
with  a  touch  of  wax,  attach  one  end  of  the  paper  to  the  stopper  of 
the  jar  ;  grease  the  stopper  with  vaseline,  and  as  quickly  as  possible 
remove  the  cork  and  replace  it  by  the  stopper  with  the  paper  sus- 
pended from  it.  Leave  it  for  some  time,  and  notice  that  the  colour 
is  not  bleached.  Afterwards  remove  the  stopper,  allow  a  drop  or 
two  of  water  to  touch  the  paper,  and  replace  it  in  the  jar.  Note  that 
the  colour  is  instantly  discharged  from  the  moistened  parts  of  the 
paper. 

This  shows  that  dry  chlorine  is  incapable  of  bleaching : 
that  the  presence  of  water  is  necessary  to  the  action.  This  is 
because  the  bleaching  action  of  chlorine  is  really  a  process  of 
oxidation.  The  chlorine  decomposes  the  water  (just  as  in  Exp. 
1 08),  liberating  oxygen,  and  uniting  with  the  hydrogen  to  form 
hydrochloric  acid  (which  we  have  learnt,  from  Exp.  103,  has 
no  bleaching  power).  Although  we  know  that  oxygen  under 
ordinary  circumstances  is  not  able  to  bleach  colours,  at  the 
moment  it  is  liberated  from  combination  it  is  endowed  with  a 
chemical  activity  which  it  has  not  got  at  other  times.  Elements 
in  this  condition  are  said  to  be  in  the  nascent  state  (that  is,  the 
newly  born  state).  Chemists  frequently  take  advantage  of  the 
extra  active  nature  of  elements  at  the  moment  of  their  libera- 
tion, to  bring  about  chemical  changes  which  they  are  incapable 
of  under  ordinary  conditions.  In  the  case  before  us,  the  nascent 
oxygen  oxidizes  the  colouring  matter,  and  converts  it  into  colour- 
less compounds. 

By  former  experiments  we  have  found  that  chlorine  does 
not  burn,  and  that  a  taper  burns  in  the  gas  with  a  very  smoky 
flame.  The  wax  of  the  taper  is  a  compound  of  hydrogen  and 
carbon,  and  from  the  manner  in  which  it  burns  in  chlorine,  we 
might  suppose  that  the  combustion  was  due  to  the  combination 
of  the  hydrogen  with  chlorine ;  and  as  the  carbon  was  set  free 
in  the  form  of  soot,  that  this  element  does  not  readily  unite 
with  chlorine.  Let  us  test  this  idea  by  a  few  experiments. 

Experiment  115. — Lower  a  burning  jet  of  coal-gas  into  chlorine. 
Notice  that  the  flame  continues  burning,  but  throws  off  a  quantity 


Combustions  in  Chlorine.  131 

of  carbon,  as  smoke,  and  also  that  fumes  are  produced  similar  to 
those  which  we  see  when  hydrochloric  acid  escapes  into  the  air. 

Experiment  116. — Boil  a  little  turpentine  in  a  test-tube,  and 
pour  some  of  it  on  a  piece  of  blotting-paper,  and  at  once  drop  it 
into  a  jar  of  chlorine.  The  turpentine  instantly  bursts  into  flame, 
again  throwing  out  a  dense  black  cloud  of  carbon,  and  forming 
fumes  of  hydrochloric  acid.  Both  coal-gas  and  turpentine  are 
compounds  of  hydrogen  and  carbon. 

Experiment  117. — Lower  a  burning  jet  of  hydrogen  into  chlorine. 
Note  that  the  flame  produces  no  black  smoke,  because  no  carbon 
is  present,  but  gives  the  same  white  fume  of  hydrochloric  acid. 

Experiment  118. — Unscrew  the  cup  from  a  deflagrating  spoon, 
and  fasten  a  piece  of  charcoal  to  the  rod  with  a  piece  of  wire. 
Make  the  charcoal  red-hot  in  a  flame  and  quickly  plunge  it  into 
chlorine.  Notice  that  the  charcoal  does  not  burn,  but  is  instantly 
extinguished.  Chlorine  will  not  combine  directly  with  carbon, 
therefore  it  is,  that  in  the  foregoing  experiments  this  element  is 
rejected  by  the  chlorine,  and  thrown  out  of  combination. 

From  these  experiments  we  see  that,  just  as  the  bleaching 
action  of  chlorine  is  the  outcome  of  its  affinity  for  hydrogen,  so 
its  peculiar  behaviour  towards  ordinary  combustibles  is  due  to 
the  same  cause.  Indeed,  these  two  elements  are  so  eager  to 
combine,  that  it  is  only  necessary  to  expose  a  mixture  of  them 
to  the  influence  of  daylight  to  bring  about  their  union.  In  dull 
daylight  the  combination  takes  place  slowly,  but  if  the  mixture 
be  exposed  for  a  single  moment  to  strong  sunshine  the  two 
gases  combine  with  an  explosion.  In  whatever  way  the  two 
elements  combine,  hydrochloric  acid  is  in  all  cases  the  product 
of  their  union.  Besides  its  powerful  affinity  for  hydrogen, 
chlorine  combines  energetically  with  most  metals.  In  many 
cases,  if  the  metal  is  in  the  form  of  powder,  or  of  thin  leaf,  the 
combination  is  so  rapid  that  the  metal  actually  takes  fire. 

Experiment  119. — Throw  into  a  jar  of  chlorine  a  small  quantity 
of  finely  divided  iron  (obtained  by  reducing  the  oxide  in  a  stream 
of  hydrogen),  or  a  little  finely  powdered  antimony.  Or  thrust  into 
the  jar  a  bundle  of  leaves  of  Dutch  metal  (that  is,  brass)  tied  to  the 
end  of  a  stout  wire.  In  each  case  the  metal  takes  fire  in  the 
chlorine,  forming  a  chloride  of  the  particular  metal. 


132  Chlorine. 

On  account  of  the  action  of  chlorine  on  metals,  this  gas 
cannot  be  collected  over  mercury,  for  it  instantly  attacks  this 
metal. 

EPITOME. 

Chlorine  is  prepared  by  acting  on  hydrochloric  acid  with 
manganese  dioxide,  or  by  acting  on  a  mixture  of  sodium  chloride 
and  sulphuric  acid  with  manganese  dioxide. 

Chlorine  is  a  greenish-yellow,  suffocating,  poisonous  gas  ;  two 
and  a  half  times  as  heavy  as  air.  It  is  moderately  soluble  in  water 
(i  litre  of  water  dissolves  about  3  litres  of  chlorine  at  the  ordinary 
temperature),  the  solution  having  the  colour  and  smell  of  the  gas. 

Chlorine  does  not  burn,  but  ordinary  combustibles  will  burn  in 
the  gas  with  a  smoky  flame. 

It  combines  readily  with  hydrogen,  producing  hydrochloric  acid, 
and  with  metals,  forming  chlorides  ;  in  many  cases  the  metal  takes 
fire  in  the  gas. 

Chlorine  is  a  powerful  bleaching  agent.  This  is  its  most 
important  property,  and  enormous  quantities  of  chlorine  are 
manufactured  by  the  action  of  manganese  dioxide  upon  hydrochloric 
acid  for  this  purpose.  As  chlorine  in  the  free  state  is  not  a  con- 
venient article  of  commerce,  the  gas  as  it  is  produced  is  made  to 
combine  with  slaked  lime,  which  gives  a  compound  known  as 
bleaching  powder.  This  substance  readily  gives  up  its  chlorine 
when  acted  on  by  acids  ;  therefore,  if  a  piece  of  coloured  material 
be  dipped  in  a  solution  of  bleaching  powder,  and  then  into  a  bath 
of  dilute  acid,  chlorine  is  liberated  in  the  pores  of  the  fibre,  and 
this  quickly  bleaches  the  colour. 

Chlorine  may  be  distinguished  from  all  other  gases  by  its  colour 
and  its  bleaching  property. 

Reactions  for  chlorine — 

(1)  From  hydrochloric  acid  and  manganese  dioxide, 

4HC1  +  MnO2  =  MnCl2  +  2H2O  +  C12 

(2)  From  salt  and  sulphuric  acid,  and  manganese  dioxide, 
2NaCl  +  2H,SO4  +  MnO2  =  Na2SO4  +  MnSO4  +  2H2O  +  Cls 

(3)  Combination  with  hydrogen,  C12  +  H2  =  2HC1 

(4)  Combination  with  metals,  (a)  Sodium,  Na  +  Cl  =  NaCl 

(£)  Copper,  Cu  +  Cl,  =.  CuCl2 
(c}  Antimony,  Sb  +  3d  =  SbCI, 


The  Halogens.  133 

Formation  of  bleaching  powder, 

CaH202  +  Cl,  =  CaOa,  +  H2O 
Action  of  acids  on  bleaching  powder, 

CaOCl2  +  H2S04  =  CaS04  +  H2O  +  Cl, 
CaOCl2  +  2HC1    =  CaCl2    +  H2O  +  C32 

The  Halogens. — Chlorine  is  one  of  a  family  of  four  elements, 
which  bear  a  close  relation  to  each  other  in  their  chemical 
behaviour,  although  they  are  not  much  alike  to  look  at. 

These  four  elements  are  fluorine  (a  gas),  chlorine  (a  gas), 
bromine  (a  liquid),  and  iodine  (a  solid).  They  are  called  the 
halogens. 

Fluorine. — This  element  is  extremely  difficult  to  obtain, 
because  it  is  so  intensely  active,  that  the  moment  it  is  set  free 
from  one  combination  it  enters  into  another,  even  combining 
with  the  materials  of  almost  any  vessel  in  which  the  experiment 
is  made. 

The  commonest  natural  compound  of  fluorine  vsfluor  spar^  that 
is,  calcium  fluoride. 

The  element  itself  is  prepared  by  passing  an  electric  current 
through  a  solution  of  potassium  fluoride  in  perfectly  pure  hydro- 
fluoric acid.  The  experiment  is  conducted  in  apparatus  constructed 
of  platinum,  which  resists  the  action  of  fluorine  better  than  any- 
thing else.  Glass  cannot  be  used  at  all,  as  even  hydrofluoric  acid 
acts  readily  on  this  substance. 

Bromine  is  a  heavy  deep-red  liquid,  which  readily  passes  off 
into  vapour  of  the  same  colour.  It  has  a  strong  suffocating 
smell  (the  name  bromine  signifies  a  stench},  and  attacks  the  nose 
and  throat,  as  chlorine  does.  It  also  irritates  the  eyes. 

In  nature  it  occurs  combined  with  sodium  and  with  magnesium 
in  the  salt  obtained  from  sea  water,  and  also  in  the  enormous  saline 
deposits  of  Stassfurt.  From  the  latter  source  the  element  is 
obtained  by  the  manufacturer. 

In  the  laboratory  we  can  prepare  bromine  from  sodium  or 
potassium  bromide,  by  mixing  either  salt  with  manganese  dioxide 
and  sulphuric  acid  (exactly  as  in  the  preparation  of  chlorine  from 
sodium  chloride).  The  sulphuric  acid  liberates  hydrobromic  acid, 
which,  in  the  presence  of  the  manganese  dioxide,  is  deprived  of  its 
hydrogen,  leaving  the  element  free. 

The  operation  is  carried  out  in  a  retort,  and  the  bromine 
collected  in  a  cooled  receiver. 


1 34  The  Halogens. 

Bromine  dissolves  in  water,  the  solution  is  called  bromine  water 
and  has  the  red  colour  of  bromine.  Like  chlorine  water,  this 
solution  of  bromine  has  bleaching  powers,  but  less  strong  than 
those  of  chlorine.  Bromine  also  combines  readily  with  metals. 

Chlorine  can  turn  bromine  out  of  its  compounds.  If  we  make 
a  solution  of  potassium  bromide  and  add  chlorine  water  to  it,  the 
chlorine  seizes  the  potassium  and  forms  potassium  chloride,  and 
the  bromine  is  set  free.  The  solution,  therefore,  turns  reddish  with 
the  liberated  bromine. 

Iodine  is  a  steel  grey  shining  crystalline  solid,  with  a  lustre 
like  a  metal.  When  gently  heated,  it  melts,  and  quickly  passes  into 
a  vapour  having  a  magnificent  violet  colour.  In  combination, 
iodine  is  present  in  small  quantities  in  sea  water,  and  is  taken  up 
by  certain  sea  plants.  Formerly  the  element  was  entirely  obtained 
from  these  sea  weeds,  but  now  it  is  chiefly  got  from  Chili  saltpetre, 
which  is  found  to  contain  a  small  quantity  of  sodium  iodate  mixed 
with  it. 

In  the  laboratory  the  element  can  be  obtained  from  potassium 
iodide,  just  as  bromine  is  from  the  bromide. 

Both  bromine  and  chlorine  can  turn  iodine  out  of  its  compounds  ; 
if,  therefore,  we  add  a  few  drops  of  either  bromine  water  or  chlorine 
water  to  a  solution  of  potassium  iodide,  the  iodine  is  liberated,  and 
the  potassium  is  taken  by  the  bromine  or  chlorine.  We  can  tell 
when  free  iodine  is  produced  by  such  reactions  as  these,  by  a 
beautiful  and  delicate  test. 

When/ra?  iodine  comes  in  contact  with  starch,  an  intensely  blue 
coloured  compound  is  formed,  therefore,  before  adding  the  bromine 
or  chlorine,  we  first  put  into  the  solution  of  potassium  iodide  a  little 
thin  starch  emulsion  (made  by  pouring  boiling  water  upon  a  little 
powdered  starch).  So  long  as  the  iodine  is  in  combination  with 
potassium,  it  is  incapable  of  uniting  with  the  starch,  but  the  instant 
the  smallest  trace  of  it  is  set  free  (by  the  addition  of  the  bromine  or 
chlorine)  the  blue  colour  is  formed. 

Compounds  of  the  halogens  with  hydrogen.— Fluorine,  bromine, 
and  iodine  each  unite  with  hydrogen,  forming  respectively  hydro- 
fluoric acid,  hydrobromic  acid,  and  hydriodic  acid.  These  three 
compounds  are  colourless  fuming  gases,  strongly  resembling  hydro- 
chloric acid  in  their  properties. 

Fluorine  combines  with  hydrogen  with  explosion,  even  in  the 
dark.  Chlorine  and  hydrogen  do  not  unite  in  the  dark,  but  do  so 
with  explosion  when  exposed  to  bright  light. 

Bromine  vapour  and  hydrogen  do  not  combine  by  the  influence 


The  Halogens.  135 

of  light,  but  do  so  in  contact  with  a  flame ;  while  iodine  and 
hydrogen  require  to  be  strongly  heated  in  order  to  cause  them  to 
unite.  This  illustrates  the  fact  that  these  four  elements  gradually 
become  less  and  less  chemically  active,  as  we  pass  from  fluorine  to 
iodine. 

Hydrofluoric  acid  has  one  remarkable  property  not  possessed 
by  either  of  the  others,  namely,  its  power  of  attacking  glass.  On 
this  account  it  is  used  to  etch  glass.  The  glass  is  first  coated  with 
wax,  and  the  design  to  be  etched  is  then  scratched  through  the 
wax.  The  whole  is  then  exposed  either  to  gaseous  hydrofluoric 
acid  or  to  a  solution  of  the  gas  in  water.  The  acid  eats  into  the 
glass  where  the  wax  has  been  removed,  and  in  this  way  etches  the 
surface.  Owing  to  its  action  on  glass,  the  acid  has  to  be  preserved 
in  either  leaden  or  guttapercha  bottles. 


CHAPTER   XIV. 

FURTHER    QUANTITATIVE   MATTERS. 

WE  have  learnt  (p.  45)  that  the  hydrogen  present  in  sulphuric 
acid  is  expelled  or  displaced  by  certain  metals,  such  as  zinc, 
magnesium,  or  iron.  Now  the  question  arises,  does  it  make 
any  difference  to  the  amount  of  hydrogen  obtained,  which  of 
these  metals  we  use  ?  In  other  words,  will  the  same  weight 
of  each  metal  displace  the  same  quantity  of  hydrogen  ?  This 
is  an  important  question,  and  we  must  try  to  answer  it  for 
ourselves  by  experiment. 


FIG.  76. 

Experiment  120.— Hydrogen  displaced  by  zinc.  Fit  up  a 
test-tube  as  shown  in  Fig.  76,  with  a  thistle  funnel  and  delivery 
tube.  [The  cork  and  tubes  must  fit  quite  tight.]  Now  weigh  out 
from  one  to  two  grams  of  zinc-foil  and  place  it  in  the  test-tube  with 


Weight  of  Hydrogen  displaced  by  Zinc          137 

a  little  water.  Support  the  apparatus  so  that  its  delivery  tube  is 
beneath  a  cylinder  standing  in  the  water  trough,  and  then  slowly 
pour  a  few  drops  of  strong  sulphuric  acid  down  the  funnel. 
Hydrogen  is  almost  immediately  set  free,  and  collects  in  the 
cylinder.  Allow  the  experiment  to  go  on  until  the  whole  of  the 
zinc  has  dissolved.  If  necessary  one  or  two  more  drops  of  acid 
can  be  added. 

When  the  action  is  finished,  the  little  apparatus  is  of  course 
filled  with  hydrogen,  but  this  is  compensated  for  by  the  air  with 
which  it  was  filled  at  the  beginning  having  been  collected  in  the 
cylinder. 

The  volume  of  gas  collected,  is  measured  as  described  in  Exp. 
100.  As  hydrogen  is  such  a  light  substance,  we  may  omit  any 
correction  for  pressure,  as  the  difference  this  would  make  will  be 
too  slight  to  effect  our  result.  The  temperature  must  be  observed 
as  before. 

EXAMPLE.  —  Weight  of  zinc  used  =  1-95  grams. 

Volume  of  hydrogen  as  collected  =  706  cc. 
Temperature  =15° 

(j)  706  cc.  of  hydrogen  at  15°,  what  volume  will  it  measure  at  o°? 


. 

(273  +  15) 

(2)  What  is  the  weight  of  670  cc.  of  hydrogen  ? 

i  litre  of  hydrogen  (1000  cc.)  weighs  0-0896  grams 
therefore  1000  :  670  :  :  0*0896  :  x 
x  =  O'o6  grams  of  hydrogen. 

Hence,  0*06  grams  of  hydrogen  are  displaced  by  1*95  grains  of  zinc  ; 
or,  i  part  by  weight  of  hydrogen  is  displaced  by  32*5  parts  by 
weight  of  zinc 

Let  us  now  repeat  this  experiment,  using  the  metal 
magnesium  instead  of  zinc.  We  can  use  the  same  apparatus 
and  conduct  the  operation  in  the  same  way,  except  that  the 
acid  must  be  more  dilute,  as  the  action  of  magnesium  is  more 
energetic  than  that  of  zinc.  We  may,  however,  perform  the 
operation  in  this  case  in  a  still  more  simple  manner. 

Experiment  121.  —  Hydrogen  displaced  by  magnesium.  Fill 
a  glass  cylinder  with  very  dilute  sulphuric  acid  (i  of  acid  to  from  20 


138  FurtJier  Quantitative  Matters. 

to  30  of  water)  and  inverc  it  in  a  glass  dish  containing  the  same 
liquid  (a  metal  pneumatic  trough  should  not  be  used).  Now 
weigh  out  about  \  gram  of  clean  magnesium  ribbon,  fold  it  up  and 
carefully  introduce  it  under  the  mouth  of  the  cylinder.  Although 
magnesium  is  heavier  than  water,  and  therefore  sinks  in  that  liquid, 
it  will  in  this  case  be  buoyed  up  to  the  top  of  the  dilute  acid  by  the 
bubbles  of  gas  which  are  evolved  from  its  surface.  In  a  very  short 
time  the  metal  will  have  entirely  dissolved,  when  the  gas  may  be 
measured  in  the  way  already  explained. 

EXAMPLE. — Weight  of  magnesium  used  =  0*48  grams 

Volume  of  hydrogen  measured  =  466  cc. 
Temperature  14°  C. 

(1)  466  cc.  of  hydrogen  at  14°  C.     What  volume  will  it  be  at  o°? 

466  x  273 
(273  +  14) 

(2)  What  is  the  weight  of  443  cc.  of  hydrogen  ? 

i  coo  cc.  hydrogen  weigh  0-0896  grams 

therefore,  1000  :  443  : :  0^0896  :  x  —  0^0397  grams  ; 

hence  0^0397  grams  of  hydrogen  are  displaced  by  0-48  grams  of 

magnesium  ; 
or,   i  part  by  weight   of  hydrogen   is  displaced   by   12  parts  of 

magnesium. 

It  is  evident  from  these  experiments  that  the  two  metals, 
zinc  and  magnesium,  differ  very  widely  in  their  power  to  dis- 
place hydrogen  from  sulphuric  acid,  for  12  parts  of  the  latter 
are  equivalent  to  32*5  parts  of  the  former  in  this  respect.  Let 
us  test  this  matter  a  little  further,  using  another  metal,  and  a 
different  hydrogen  compound. 

Experiment  1 22.— Hydrogen  displaced  by  sodium.  Take  one 
of  the  little  lead  tubes  described  on  p.  42,  and  counterpoise  it  on  the 
balance.  Then  fill  the  tube  with  sodium,  as  there  explained,  using 
a  clean  piece  and  not  exposing  it  to  the  air  longer  than  is  absolutely 
necessary.  Then  weigh  the  tube  and  contents.  Now  take  the  tube 
between  the  thumb  and  forefinger  and  quickly  introduce  it  beneath 
the  mouth  of  a  cylinder  of  water  in  the  pneumatic  trough,  holding 
it  there  until  no  more  gas  is  evolved.  When  the  gas  stops,  gently 
shake  the  little  tube  by  tapping  it  against  the  bottom  of  the  trough 
so  as  to  be  quite  sure  that  the  sodium  is  entirely  gone.  Measure 
the  gas,  and  take  its  temperature  as  before. 


Equivalent   Weights  of  Elements.  139 

EXAMPLE. — Weight  of  sodium  used  =  0*3.1.  grams. 

Volume  of  hydrogen  as  collected  =173  cc. 
Temperature  =  12°  C. 

(1)  173  cc.  at  12°,  what  volume  at  o°? 

173^7.3^6 

273  +  12 

(2)  What  is  the  weight  of  165  cc.  of  hydrogen  ?    V 

loco  cc.  hydrogen  weigh  0*0896  grams  ; 

.*.  loco  :  165  : :  0*0896  :  x  —  0*0147  grams. 
Hence  0*0147  grams  of  hydrogen  are  displaced  by  0*34  grams  of 

sodium ; 
or,  i  part  by  weight  of  hydrogen  is  displaced  by  23  parts  of  sodium. 

This  experiment  teaches  us  that  sodium  is  different  from 
both  zinc  and  magnesium  in  its  power  of  turning  hydrogen 
out  of  its  compounds,  for  we  find  that  in  this  respect  23  parts 
of  sodium  are  equivalent  to  12  parts  of  magnesium,  and  to 
32*5  parts  of  zinc. 

In  other  words,  we  may  say  that  23  parts  of  sodium,  12 
parts  of  magnesium,  and  32-5  parts  by  weight  of  zinc,  are  each 
chemically  equivalent  to  i  part  by  weight  of  hydrogen. 

Every  element,  however,  is  not  capable  of  displacing 
hydrogen  from  its  compounds  in  this  simple  way,  or  it  would 
be  an  easy  matter  to  find  the  particular  weight  of  each  element 
which  was  equivalent  to  i  part  of  hydrogen. 

In  the  case,  therefore,  of  a  great  many  of  the  elements, 
we  have  to  go  to  work  in  a  more  roundabout  manner  in 
order  to  find  what  weight  of  them  is  equivalent  to  i  part 
by  weight  of  hydrogen.  For  instance,  having  discovered  the 
weight  of  zinc  which  is  capable  of  displacing  one  part  by 
weight  of  hydrogen,  we  can  find  the  weight  of  certain  other 
metals  which  can  be  displaced  by  this  equivalent  weight  of 
zinc. 

Experiment  123. — Add  a  few  drops  of  a  solution  of  copper 
sulphate  to  half  a  test-tube  full  of  water,  and  put  into  the  solution 
a  narrow  strip  of  sheet  zinc.  Notice  that  the  zinc  at  once  begins 
to  blacken,  and  that  in  a  very  short  time  the  blue  colour  of  the 
solution  begins  to  fade  away,  and  presently  to  disappear  altogether. 


140  Further  Quantitative  Matters. 

Experiment  124.— Pour  two  or  three 'drops  of  silver  nitrate 
solution  upon  a  piece  of  flat  glass,  and  lay  a  minute  piece  of  zinc 
upon  the  liquid,  and  let  it  remain  still.  Look  at  it  with  a  pocket 
lens,  and  notice  bright  shining  crystals  of  silver  beginning  to  be 
deposited  all  round  the  piece  of  zinc,  and  gradually  spreading  right 
across  the  drops  of  liquid. 

These  two  experiments  show  us  that  both  copper  and 
silver  are  displaced  from  solutions  of  their  salts  by  the  metal 
zinc,  for  the  black  deposit  in  the  first  case  is  simply  copper  in 
the  form  of  a  very  fine  powder.  We  might  repeat  these 
experiments,  employing  magnesium  in  the  place  of  the  zinc, 
with  similar  results.  In  this  case,  however,  the  action  is 
slower,  and  the  displacement  of  the  metals  is  less  complete. 

In  order  to  ascertain  the  weight  of,  say,  silver  which  is 
displaced  by  32*5  parts  by  weight  of  zinc,  we  can  proceed  as 
follows. 

Experiment 125. — Silver  displaced  by  zinc.  Weigh  out  about 
half  a  gram  of  clean  zinc  foil,  place  it  in  a  small  beaker  and  pour 
a  solution  of  silver  nitrate  upon  it,  and  allow  it  to  stand.  As 
the  silver  is  gradually  deposited,  the  zinc  is  as  gradually  dissolved, 
until  finally  it  has  entirely  disappeared.  By  gently  feeling  with  a 
glass  rod  it  is  easy  to  ascertain  when  the  whole  of  the  zinc  is 
dissolved,  because  the  precipitated  silver  is  quite  soft  and  spongy. 
[The  solution  of  silver  nitrate  should  not  be  too  dilute,  or  it  may 
be  necessary  to  decant  off  the  first  quantity  and  add  more,  in 
order  to  provide  enough  silver  for  the  weight  of  zinc  employed.] 
When  the  zinc  is  completely  gone,  the  liquid  must  be  carefully 
decanted  off,  and  the  precipitated  silver  washed  once  or  twice  by 
half  filling  the  beaker  with  water,  stirring  gently,  allowing  it  to 
settle  and  again  decanting  the  liquid.  The  silver  is  next  transferred 
to  a  clean  porcelain  crucible,  which,  with  its  lid,  has  been  counter- 
poised carefully.  This  is  done  by  collecting  it  together  as  much 
as  possible,  and  withdrawing  it  by  means  of  a  glass  rod,  in  the 
manner  shown  in  Fig.  77,  the  last  particles  being  rinsed  out 
with  water.  When  the  whole  has  in  this  way  been  got  into  the 
crucible,  the  water  is  decanted  and  the  silver  drained  as  dry  as 
possible.  The  crucible,  partly  covered  with  its  lid,  is  then  placed 
in  a  hot-air  oven,  heated  to  about  110°  to  120°,  in  order  to  dry  the 
silver  perfectly  ;  after  which  it  is  allowed  to  cool,  and  weighed. 
It  should  then  be  returned  to  the  oven  for  a  second  heating,  and 


Equivalent   Weights  of  Elements. 


141 


once  more  weighed.     If  this  last  weighing  agrees  with  the  one 
before,  it  shows  that  the  substance  was  perfectly  dry. 
EXAMPLE. — Weight  of  zinc  used  =  0-48  gram. 

„      silver  obtained  =  r6o  grams. 
Then  as  ©'48  :  32^5  : :  1*6  :  x. 


x  = 


_  32*5  x  r6  _ 


=  io8'3  =  Weight  of  silver  displaced  by  32*5  parts 


0-48 
by  weight  of  zinc. 

By  this  experiment,  then,  108-3  Parts  by  weight  of  silver 
are  equivalent  to  3  2 '5  parts  by  weight  of  zinc.  But  this 
weight  of  zinc  is  equivalent  to  i  part  by  weight  of  hydrogen  ; 


FIG.  77. 

therefore  we   say    that   108*3    Parts   by  weight  of  silver  are 
equivalent  to  i  part  by  weight  of  hydrogen. 

By  a  similar  experiment,  using  a  solution  of  copper  sulphate 
instead  of  silver  nitrate,  we  could  ascertain  the  weight  of  copper 
that  is  equivalent  to  i  part  of  hydrogen. 

When  we  examine  the  compound  of  hydrogen  with 
chlorine,  we  find  that  the  elements  combine  together  in  the 
proportion  of  one  part  by  weight  of  hydrogen  with  35-5  parts 
by  weight  of  chlorine.  We  say,  therefore,  that  35*5  parts  of 
chlorine  are  equivalent  to  i  part  of  hydrogen. 

By  Exp.  122  we  learnt  that  23  parts  of  sodium  are  equiva- 
lent to  i  part  of  hydrogen.     Hence  we  have — 
35*5  parts  by  weight  of  chlorine  equivalent  to  i  part  of  hydrogen. 
23         „  „  sodium  „  i  „ 

Now   we    know    that    sodium    and    chlorine    themselves 


142  Further  Quantitative  Matters, 

combine  together,  and  the  question  arises,  do  they  unite  in 
the  proportion  of  these  equivalent  numbers  ? 

When  the  compound  sodium  chloride  is  analyzed,  it  is 
found  to  contain  the  two  elements  in  the  proportion  i  part  of 
sodium  to  1*535  parts  of  chlorine, 

but  i  :  1-535  •:  23  :35'5; 

therefore    the    proportion    in    which    sodium    and    chlorine 
separately  combine   with  hydrogen,   is    the    same  as  that    in 
which  they  unite  together.     We  may  say,  therefore,  that  23 
parts  of  sodium  are  equivalent  to  35^5  parts  of  chlorine. 
Again,  from  Exp.  97  it  was  found  that — 

i  part  of  hydrogen  combines  with  8  parts  of  oxygen  by  weight, 
and  from  Exp.  94  that — 

1 2  parts  of  magnesium  combine  with  8  parts  of  oxygen  by 
weight.1 

But  we  have  also  learnt  that  12  parts  of  magnesium  and 
i  part  of  hydrogen  are  chemically  equivalent,  therefore  they 
are  each  equivalent  to  8  parts  of  oxygen. 

Similarly  since  23  parts  of  sodium  and  8  parts  of  oxygen 
are  each  equivalent  to  i  part  of  hydrogen,  they  are  equivalent 
to  each  other,  and  if  we  were  to  examine  the  compound  they 
form  when  they  unite,  we  should  find  that  it  contained  sodium 
and  oxygen  in  the  proportion  of  23  to  8. 

And  again,  since  35^5  parts  of  chlorine  and  12  parts  of 
magnesium  are  each  equivalent  to  i  part  of  hydrogen,  they 
are  equivalent  to  each  other;  and  when  the  compound  of 
magnesium  and  chlorine  is  analyzed  it  is  actually  found  to 
contain  these  elements  in  the  proportion  of  12  parts  of 
magnesium  to  35*5  of  chlorine. 

By  Exp.  96  we  have  learnt  that  sulphur  and  oxygen  combine 
in  equal  weights,  that  is — 

8  parts  of  sulphur  combine  with  8  parts  of  oxygen ; 
but 

i  part  of  hydrogen  combines  with  8  parts  of  oxygen. 

1  The  proportion  i  '5  magnesium  to  i  oxygen,  obtained  by  Exp,  94,  is 
the  same  as  12  to  8. 


Equivalent    Weights  of  Elements.  143 

Do  sulphur  and  hydrogen  follow  the  same  rule  as  in  the 
former  cases,  and  combine  in  the  proportion  of  8  to  i  ?  When 
the  compound  of  sulphur  and  hydrogen  is  examined,  it  is 
found  that  it  contains  the  elements  in  the  proportion,  sulphur 
1 6  parts  to  hydrogen  i  part.  Instead  of  8  to  i,  the  pro- 
portion, therefore,  is  8  X  2  of  sulphur  to  i  part  of  hydrogen. 
When  chemists  had  investigated  a  large  number  of  cases 
where  elements  which  combined  together,  also  combined 
separately  with  some  other  element,  they  found  that  in  some 
instances,  the  proportion  in  which  the  elements  separately 
combined  with  one  common  element,  was  the  same  as  that 
in  which  they  united  together ;  while  in  other  cases  (as  in 
the  above  example)  it  is  a  simple  multiple  of  that  number. 
These  facts  are  expressed  in  the  third  general  law  of  chemical 
combination,  known  as  the 

Law  of  reciprocal  proportions,  or  the  law  of  equivalent  pro- 
portions. The  weights  of  different  elements  which  separately 
combine  with  a  constant  weight  of  another  element,  are  either  the 
same  as,  or  are  simpler  multiples  of  the  weights  of  these  elements 
which  combine  with  each  other. 

These  various  weights  of  the  elements  are  called  their 
combining  proportions,  or  the  equivalent  weights  of  the  elements. 
They  represent  the  proportions  by  weight  in  which  the 
elements  are  able  to  combine  together,  relative  to  i  part  by 
weight  of  hydrogen. 

Experiments  on  Neutralizing  Acids. 

Experiment  126. — Counterpoise  a  small  beaker,  and  then  weigh 
into  it  100  grams  of  strong  oil  of  vitriol.  Pour  this  into  a  litre 
flask,  rinsing  the  beaker  three  or  four  times  with  water,  and  care- 
fully pouring  each  rinsing  into  the  flask,  so  that  the  whole  of  the 
acid  may  be  transferred  without  loss.  Now  fill  the  flask  up  to 
the  mark  with  distilled  water.  This  litre,  therefore,  contains  100 
grams  of  sulphuric  acid  ;  *  and  as  a  litre  is  1000  cc.,  each  cc.  of 
the  liquid  will  contain  0*1  gram  of  acid.  Transfer  this  solution 
to  a  clean  stoppered  bottle.  Next  weigh  out  25  grams  of  solid 
caustic  soda.  [This  weighing  must  be  done  as  quickly  as  possible, 

1  In  reality  this  is  not  quite  exact,  as  even  the  strong  sulphuric  acid 
contains  a  small  quantity  of  water  ;  but  it  will  be  near  enough  for  our 
present  purpose. 


144  Further  Quantitative  Matters. 

as  caustic  soda  absorbs  moisture  from  the  air  very  rapidly.]  Dis- 
solve this  in  water  in  a  half-litre  flask,  and  then  fill  it  to  the  mark 
with  more  water.  Since  the  half  litre  (that  is,  500  cc.)  contains 
25  grams  of  caustic  soda,  i  cc.  will  therefore  contain  0-05  gram. 

By  means  of  a  pipette  (Fig.  65),  transfer  50  cc.  of  this  solution 
to  a  beaker,  and  add  to  it  one  or  two  drops  of  litmus  solution. 

Fill  a  burette  (Fig.  64)  with  the  prepared  sulphuric  acid,  and 
allow  it  to  run,  in  small  quantities  at  a  time,  into  the  beaker,  which 
should  be  stood  on  a  white  tile,  or  piece  of  white  paper.  As  the 
point  is  approached  at  which  the  alkali  is  neutralized,  each  drop 
of  acid  produces  a  temporary  reddening  of  the  liquid  as  it  falls  in. 
The  liquid  should  be  stirred,  and  the  addition  of  acid  stopped  the 
instant  the  blue  litmus  is  permanently  reddened.  Read  and  note 
how  many  cc.  of  acid  have  been  added  in  order  to  so  neutralize 
the  alkali,  and  then  repeat  the  experiment  so  as  to  obtain  two 
results  which  agree. 

EXAMPLE. — 50  cc.  of  the  caustic  soda  required  30  cc.  of  the 
sulphuric  acid.  Since  I  cc.  of  the  acid  contains  o'l  gram  of  sul- 
phuric acid,  30  cc.  will  contain  3  grams,  therefore  3  grams  of 
sulphuric  acid  are  neutralized  by  50  cc.  of  the  caustic  soda,  or  2'5 
grams  ;  or  100  grams  of  acid  are  neutralized  by  83*3  grams  of 
caustic  soda. 

Experiment  127. — Weigh  out  25  grams  of  caustic  potash,  and 
dissolve  it  in  a  half-litre  flask,  and  fill  to  the  mark  with  water. 
Take  50  cc.  of  this  solution  and  proceed  exactly  as  in  the  last 
experiment.  Note  the  volume  of  acid  required  to  exactly  neutralize 
the  alkali. 

EXAMPLE. — 50  cc.  of  caustic  potash  solution  required  22  cc. 
of  acid,  or  2'2  grams.  Therefore  2*2  grams  of  sulphuric  acid  are 
neutralized  by  2*5  grams  of  caustic  potash ;  or  100  grams  of  acid 
are  neutralized  by  1 13*6  grams  of  caustic  potash. 

According  to  these  experiments,  therefore,  83-3  parts  by 
weight  of  caustic  soda  are  equivalent,  in  their  power  of 
neutralizing  sulphuric  acid,  to  113-6  parts  of  caustic  potash. 

Let  us  now  see  whether  these  two  alkalies  stand  in  the 
same  relation  towards  another  acid. 

Experiment  128.— Take  50  cc.  of  ordinary  concentrated  hydro- 
chloric acid,  by  means  of  a  pipette,  and  dilute  it  to  half  a  litre. 
(For  our  present  purpose  it  is  not  necessary  to  know  the  weight 
of  the  acid  we  arc  going  to  use.) 


Acids  and  Alkalies.  145 

Now  withdraw  50  cc.  of  this  diluted  acid  with  a  pipette,  transfer 
it  to  a  small  beaker  and  colour  it  with  litmus. 

Fill  a  burette  with  the  caustic  soda  solution  prepared  in  Exp. 
1 26,  and  add  it  gradually  to  the  acid  until  the  solution  is  exactly 
neutral. 

Now  repeat  this  with  the  caustic  potash  solution,  taking  another 
50  cc.  of  the  acid  and  neutralizing  it  with  the  potash. 

Calculate  as  in  the  former  experiment  the  different  weights  of 
these  two  alkalies  required  to  neutralize  equal  quantities  of  hydro- 
chloric acid  ;  and  see  whether  they  are  in  the  same  ratio  as  83-3 
to  1136. 


CHAPTER   XV. 

THE   ATOMIC    THEORY. 

THE  three  laws  of  chemical  combination,  namely,  the  law  of 
constant  proportion,  the  law  of  multiple  proportion,  and  the 
law  of  equivalent  proportions,  are  the  general  expressions  of 
a  vast  number  of  well-established  facts.  By  experiments 
such  as  those  described  in  the  foregoing  chapters,  chemists 
have  discovered  the  facts  that  chemical  combination  between 
the  elements  takes  place  according  to  these  rules  or  laws,  facts 
which  must  be  regarded  as  indisputable. 

The  human  mind,  however,  is  not  satisfied  with  facts  •  it 
always  asks  the  question,  Why  ?  In  order  to  account  for,  and 
to  explain  these  extraordinary  facts,  in  order  to  give  some 
satisfactory  answer  to  the  question,  Why  do  the  elements  unite 
in  definite  and  in  multiple  proportions  ?  chemists  adopt  a  theory. 
A  law  is  the  general  expression  of  discovered/^/.?  /  a  theory 
is  a  hypothesis  or  guess  which  we  make  in  order  to  explain 
the  facts.  If  the  hypothesis  fits  all  the  facts,  it  is  then  called 
a  theory ;  and  so  long  as  no  new  facts  are  discovered  which 
do  not  square  with  the  theory,  it  is  accepted  as  an  explanation 
of  the  facts ;  but  if  fresh  facts  are  found  out  which  will  not 
harmonize  with  the  theory,  then  the  theory  must  be  given  up, 
and  a  new  one  proposed. 

The  theory  which  chemists  have  adopted  to  explain  the 
facts  connected  with  chemical  combination,  is  known  as  the 
atomic  theory,  or  sometimes  Daltoris  atomic  theory,  from 
the  name  of  the  chemist  who  first  proposed  it. 

Dalton's  theory  consists  of  four  hypotheses  or  surmises  : — 

(i)  That  all  the  elements,  whether  solid,  liquid,  or  gaseous, 
consist  of  a  vast  number  of  minute  indivisible  particles  or  atoms. 


The  Atomic  Theory.  147 

(2)  That  the  atoms  of  any  one  element  have  all  the  same 
weight. 

(3)  That   the  atoms  of  different  elements  have  different 
weights,  and  that  these  weights  stand  in  the  same  ratio  to 
one  another  as  the  numbers  expressing  the  combining  pro- 
portions of  the  elements. 

(4)  That    chemical   combination   consists   simply   in   the 
union  together  of  atoms ;  the  atoms  being  held  together  by 
the  operation  of  the  force,  chemical  affinity. 

Let  us  see  now  how  this  theory  will  help  us  to  understand 
the  known  facts  about  the  manner  in  which  elements  combine 
together  ;  and  first  as  to  the  law  of  constant  proportion.  When 
sodium  combines  with  chlorine,  we  have  seen  (p.  142)  that  the 
proportion  in  which  these  two  elements  unite,  is  sodium  23 
parts  and  chlorine  35^5  parts.  Analysis  shows  that  this  pro- 
portion is  constant,  that  the  compound  sodium  chloride  in- 
variably contains  the  elements,  sodium  and  chlorine,  in  these 
proportions, 

According  to  the  atomic  theory,  sodium  chloride  is  the 
product  of  the  union  of  atoms  of  sodium  with  atoms  of  chlorine, 
and  the  relative  weights  of  these  atoms  is  expressed  by  the 
combining  proportions  of  the  elements,  namely,  23  and  35*5. 
Since,  by  the  theory,  atoms  are  indivisible,  it  follows  that  the 
compound  produced  by  the  union  of  one  atom  of  each  ot 
these  two  elements  must  always  have  the  same  composition. 
The  proportion  cannot  be  22  :  35-5,  or  23  :  30,  because  neither 
of  these  represents  the  relative  weights  of  the  indivisible  atoms 
of  sodium  and  chlorine. 

Again,  as  to  the  law  of  multiple  proportions.  We  saw 
(p.  1 1 8)  that  carbon  unites  with  oxygen  in  two  proportions, 
producing  two  compounds. 

In  the  first  of  them  the  elements  are  present  in  the  pro- 
portion, carbon  i  part,  and  oxygen  1-33  parts.  From  the 
figures  on  p.  142,  it  will  be  seen  that  the  combining  proportion 
for  oxygen  is  8.  But  the  weight  of  carbon  which  bears  the 
same  ratio  to  8  as  i  bears  to  1*334  is  6. 

i  :  1-334::  6  :8 


148  The  Atomic  Theory. 

Therefore  carbon  and  oxygen  are  contained  in  this  com- 
pound in  the  proportion  of  carbon  6  parts  and  oxygen  8  parts. 
The  composition  of  this  compound,  like  that  of  the  sodium 
chloride,  is  invariable ;  one  indivisible  atom  of  carbon  unites 
with  one  indivisible  atom  of  oxygen,  the  relative  weights  of 
these  two  atoms  being  as  6  :  8. 

In  the  second  compound  of  carbon  and  oxygen,  the  carbon 
and  oxygen  are  present  in  the  proportion,  carbon  6  parts 
and  oxygen  16  parts.  The  atomic  theory  accounts  for  this. 
The  one  atom  of  carbon  is  here  combined  with  two  atoms  of 
oxygen,  each  having  the  relative  weight  of  8.  If  carbon  is  to 
combine  with  more  than  i  atom  of  oxygen,  since  atoms  are 
indivisible,  it  must  unite  with  at  least  2  atoms. 

In  the  five  compounds  of  nitrogen  and  oxygen  (p.  118)  the 
two  elements  are  present  in  the  proportions — 


(1)  Nitrogen 

(2)  Nitrogen 

(3)  Nitrogen 

(4)  Nitrogen 

(5)  Nitrogen 


oxygen  —  i 
oxygen  —  i 
oxygen  —  I 
oxygen  =  i 
oxygen  =  I 


o'S7  or  14  :  8 
1*14  or  14:  16 
171  or  14  :  24 
2-28  or  14  :  32 
2-85  or  14  140 


The  increase  in  oxygen  takes  place  regularly  by  the  addition 
of  a  weight  of  that  element  equal  to  8,  that  is,  by  a  weight 
equal  to  the  combining  proportion  of  oxygen,  which  according 
to  Dalton's  hypothesis  is  the  relative  weight  of  the  atom  of 
oxygen. 

Atomic  Weights  of  the  Elements.— As  an  outcome  of 
the  atomic  theory,  the  numbers  which  formerly  represented 
the  combining  proportions  of  the  elements  became  invested 
with  a  new  significance,  and  were  called  the  atomic  weights  of 
the  elements.  Since  Dalton's  day,  however,  with  the  growth 
of  chemical  knowledge,  it  has  been  found  necessary  to  change 
many  of  the  numbers  which  were  by  him  regarded  as  the 
atomic  weights;  and  in  many  cases,  instead  of  the  number 
which  expresses  the  combining  proportion,  it  is  some  multiple 
of  this  number  which  is  to-day  accepted  as  the  true  atomic 
weight.  For  example,  the  combining  proportion  of  oxygen 
is  8,  its  atomic  weight  is  now  regarded  as  16.  The  combining 
proportions  of  maenesium  and  of  sulphur  are  respectively  12 


Atoms  and  Molecules.  149 

and  1 6,  but  we  now  assign  the  numbers  24  and  32  as  the 
atomic  weights  of  these  elements.1 

The  numbers  which  are  at  the  present  time  accepted  by 
chemists  as  the  approximate  atomic  weights  of  the  elements, 
are  given  in  the  third  column  of  the  table  on  the  cover.  The 
student  should  make  himself  familiar  with  the  atomic  weights 
of  a  number  of  the  commoner  elements.  Referring  to  the 
table,  we  there  see  that  the  atomic  weight  of  carbon  is  1 2  ;  let 
him  remember  that  this  simply  means,  that  the  smallest  weight 
of  carbon  which  enters  into  chemical  combination  is  12  times 
as  heavy  as  the  smallest  weight  of  hydrogen  which  is  capable 
of  uniting  with  another  element ;  or,  in  the  words  of  the  atomic 
theory,  that  the  carbon  atom  is  12  times  as  heavy  as  the 
hydrogen  atom.  The  atomic  weight  of  oxygen  and  of 
nitrogen  are  16  and  14,  that  is  to  say,  the  atoms  of  these 
elements  are  respectively  16  and  14  times  heavier  than  the 
hydrogen  atom.  Hydrogen  being  the  lightest  known  substance, 
the  weight  of  its  atom  is  taken  as  the  unit,  and  in  the  table 
therefore  its  atomic  weight  is  given  as  i. 

Atoms  and  Molecules. — Chemists  believe  that,  generally 
speaking,  atoms  are  unable  to  exist  alone,  that  is,  in  a  free  or 
uncombined  state ;  but  that  the  moment  an  atom  is  expelled 
from  combination,  it  immediately  combines  with  some  other 
atom  or  atoms ;  either  with  an  atom  like  itself  or  an  atom  of 
some  other  element.  The  smallest  particle  of  matter,  there- 
fore, which  is  capable  of  independent  existence,  consists 
usually  of  a  little  group  or  system  of  atoms,  bound  together 
by  chemical  affinity.  These  groups  of  atoms  are  called 
molecules.  If  the  atoms  which  are  thus  associated  together  to 
form  a  molecule  are  atoms  of  different  elements,  the  molecule 
is  the  molecule  of  a  compound,  but  when  an  atom  is  combined 
only  with  other  atoms  of  the  same  element,  the  molecule  is 
a  molecule  of  an  element.  For  example,  a  molecule  of  the 

1  There  are  many  considerations  which  now  influence  chemists,  in 
deciding  upon  the  particular  number  which  shall  be  assigned  as  the  atomic 
weight  of  an  element,  but  the  study  of  these  matters  the  student  may  leave 
until  a  later  stage,  and  he  will  find  them  fully  set  forth  in  more  advanced 
text-books.  See  Newih's  "  Inorganic  Chemistry." 


150  The  Atomic  Theory. 

compound  sodium  chloride  consists  of  one  atom  of  sodium 
united  to  one  atom  of  chlorine,  while  a  molecule  of  the 
element  chlorine  is  composed  of  an  atom  of  chlorine  united  to 
another  chlorine  atom.  To  the  mind  of  the  chemist  all  kinds 
of  matter,  say  a  piece  of  chalk,  or  a  drop  of  water,  present 
the  appearance  of  an  innumerable  multitude  of  these  tiny 
molecules1  or  groups  of  atoms,  packed  more  or  less  closely 
together.  When  a  solid  substance,  such  as  ice,  is  melted,  the 
mental  picture  which  the  chemist  sees  of  the  operation,  is  that 
the  molecules  are  thrust  further  apart  from  each  other,  and 
again  when  the  liquid  water  is  converted  into  steam,  he  sees 
the  molecules  separated  to  a  still  greater  distance  from  each 
other.  But  in  all  these  changes,  he  sees  that  the  atoms  which 
compose  the  molecules  still  remain  associated  together;  each 
water  molecule  is  a  little  group  of  three  atoms  (two  of  hydrogen 
and  one  of  oxygen),  and  these  three  atoms,  held  together  by 
chemical  affinity,  remain  closely  bound  together  whether  the 
water  be  solidified  to  ice  or  vaporized  to  steam.  These 
changes  only  effect  the  molecule  as  a  whole ,  they  are,  there- 
fore, only  physical  changes;  but  as  soon  as  we  disturb  the 
composition  of  the  molecule,  as  soon  as  we  break  through  the 
force  which  unites  the  atoms  together  in  the  group,  then  we 
produce  a  chemical  change  upon  the  substance.  We  then 
separate  the  constituent  atoms  in  the  water  molecules,  and 
obtain  two  substances,  namely,  hydrogen  and  oxygen,  which 
are  both  entirely  different  in  properties  from  the  original  water. 
We  may  see  some  analogy  to  the  relations  between  atoms 
and  molecules,  in  many  of  the  systems  of  the  heavenly  bodies. 
Thus,  the  planet  Jupiter  with  its  five  moons  may  be  compared 
to  a  molecule  consisting  of  six  atoms.  The  planet  and  its 
satellites  remain  constantly  associated  together,  performing 
certain  movements  relative  to  each  other,  while,  at  the  same 
time,  the  whole  system  in  its  undisturbed  unity  courses  through 
space  on  its  own  independent  orbit.  This  illustration,  how- 
ever, is  incomplete  in  this  point;  we  know  that  the  force, 

1  The  word  "molecule"  means  little  mass.  They  are  so  extremely 
minute  that  a  single  drop  of  water  is  made  up  of  countless  millions.  It 
is  utterly  beyond  the  power  of  the  strongest  microscope  to  detect  them. 


Atoms  and  Molecules.  15! 

which  binds  the  planet  and  its  moons  together  and  keeps  them 
united  in  a  single  system,  is  the  same  as  that  which  regulates 
the  movements  of  the  united  system  as  it  travels  round  the 
central  sun,  namely,  gravitation  ;  but  the  force  which  binds  the 
atoms  together  in  a  molecule,  is  the  force  we  call  chemical 
affinity,  and  it  appears  to  be  a  totally  different  force  from  any 
of  the  merely  physical  forces  which  operate  between  the 
molecules  themselves. 

An  atom  may  be  defined  as  the  smallest  weight  of  matter 
which  can  take  part  in  a  chemical  change:  while  we  define  a 
molecule  as  the  smallest  weight  of  matter  which  can  exist  in  the 
free  state. 


CHAPTER  XVI. 

CHEMICAL   SYMBOLS,    FORMULAE   AND    EQUATIONS. 

Chemical  Symbols. — In  the  second  column  of  the  table 
on  the  cover,  the  symbols  are  given  which  chemists  use  to 
denote  the  different  elements.  In  a  number  of  cases  these 
symbols  are  merely  the  first  letter  of  the  name,  such  as  Carbon, 
C ;  Hydrogen,  H  ;  Oxygen,  O  ;  Sulphur,  S ;  and  so  on.  In 
others  it  is  either  the  first  or  the  second  letters,  as  Bromine, 
Br;  Silicon,  Si;  or,  the  first  and  one  other  that  is  prominent 
in  the  name,  such  as  Chlorine,  Cl  \  Manganese,  Mn.  In  some 
instances  the  symbol  is  taken  from  the  Latin  name,  thus — 
Copper  (cuprum],  Cu ;  Iron  (ferrum),  Fe;  Silver  (argentum], 
Ag.1  These  symbols  are  not  intended  to  be  used  as  mere 
shorthand  signs  only,  but  they  stand  in  all  cases  for  one  atom 
of  the  various  elements.  Thus,  the  symbol  H  stands  for  one 
atom  of  hydrogen ;  the  symbols  O,  K,  Na,  Cl,  represent  one 
atom  of  oxygen,  potassium,  sodium,  and  chlorine  respectively, 
and  not  any  indefinite  quantity  of  these  substances.  They 
are,  in  a  strict  sense,  atomic  symbols,  that  is,  symbols  of  the 
atoms,  and  they  should  therefore  never  be  employed  by  the 
student  as  mere  abbreviations  for  the  names  of  the  elements.2 
If  we  know  how  many  atoms  of  an  element  go  to  make  up 

1  The  student  must  make  himself  familiar  with  the  symbols  of  all  the 
more  common  elements,  such  as  those  contained  in  the  list  on  p.  8. 

2  It  is  true  that  chemists  often  fall  into  the  habit  of  employing  these 
symbols  as  mere  abbreviations  in  writing,  but  the  habit  is  to  be  deprecated, 
especially  in  the  beginner,  as  its  practice  tends  greatly  to  obscure  the  true 
significance  of  the  symbols  ;  namely,  that  of  representing  a  definite  quantity 
(one  atom)  of  the  various  elements. 


Chemical  Formula.  153 

the  molecule  of  that  element,  we  can  express  that  knowledge 
by  means  of  these  symbols,  by  the  use  of  numerals  placed 
immediately  after  them.  Thus,  in  the  case  of  the  elements 
hydrogen,  oxygen,  nitrogen,  chlorine,  and  some  others,  the 
molecules  consist  of  groups  of  two  atoms,  and  as  the  symbol 
stands  for  one  atom,  the  molecules  of  these  elements  are 
represented  by  H^  O2,  N2,  C12,  respectively.  The  symbol 
H2,  therefore,  means  one  molecule  of  hydrogen. 

Chemical  Formulae* — By  means  of  these  symbols,  the 
composition  of  compounds  is  represented.  This  is  accom- 
plished by  placing  the  symbols  for  the  various  elements  in  the 
compound,  side  by  side.  Thus,  the  compound  obtained  when 
sodium  (symbol  Na)  and  chlorine  (symbol  Cl)  combine  to- 
gether, is  expressed  by  the  united  symbols  of  the  two  elements, 
NaCl.  Such  an  arrangement  of  symbols  standing  for  a  com- 
pound is  called  the  formula  for  that  compound.  Such  a 
formula,  however,  expresses  the  composition  of  a  molecule  of  the 
compound ;  and  as  Na  stands  for  one  atom  of  sodium,  and  Cl 
for  one  atom  of  chlorine,  the  molecule  of  sodium  chloride  (as 
represented  by  this  formula)  will  consist  of  one  atom  of  each 
of  the  two  constituents.  HC1,  in  like  manner,  is  the  molecular 
formula  for  the  compound  of  hydrogen  and  chlorine,  and  the 
formula  states  that  the  molecule  of  hydrochloric  acid  consists 
of  one  atom  of  hydrogen  united  to  one  atom  of  chlorine. 

When  the  molecules  of  any  compound  contain  more  than 
one  atom  of  any  particular  element,  the  fact  is  indicated  by  the 
use  of  small  numbers  placed  immediately  after  the  symbol  of 
that  element ;  thus  H2O  is  the  molecular  formula  for  water,  a 
compound  containing  in  its  molecule  two  atoms  of  hydrogen 
and  one  of  oxygen.  NH4C1  is  the  formula  for  one  molecule 
of  ammonium  chloride,  in  which  the  molecule  is  composed  of 
one  atom  of  nitrogen,  four  atoms  of  hydrogen,  and  one  of 
chlorine.  Very  often  in  formulae  a  little  more  complex,  it  is 
necessary  to  indicate  the  presence  of  certain  groupings  of 
atoms.  For  this  purpose  brackets  are  employed,  thus 
(NH4)2SO4  represents  a  molecule  containing  four  atoms  of 
oxygen,  one  atom  of  sulphur,  eight  atoms  of  hydrogen,  and 
two  of  nitrogen,  the  last  two  elements  being  associated  together 


154     Chemical  Symbols,  Formula  and  Equations. 

in  two  groups,  each  group  consisting  of  one  atom  of  nitrogen 
and  four  of  hydrogen. 

When  it  is  necessary  to  indicate  more  than  one  molecule 
of  any  substance,  large  numerals  are  placed  before  the  symbol 
or  formula,  thus  2H2  means  two  molecules  of  hydrogen  ;  5H2O 
signifies  five  molecules  of  water. 

Chemical  Equations. — One  of  the  chief  uses  of  these 
symbols  and  formulae  is  to  enable  chemists  to  express,  in  a 
condensed  and  precise  form,  a  considerable  amount  of  infor- 
mation respecting  the  various  chemical  changes  which  it  is 
their  business  to  study.  These  changes  are  called  chemical 
reactions,  and  they  are  expressed  in  the  form  of  equations.  The 
symbols  and  formulae  of  all  the  materials  undergoing  change 
are  placed  to  the  left,  and  those  of  the  new  products  'resulting 
from  the  action,  on  the  right. 

Thus,  when  mercury  and  iodine  unite  to  form  mercuric 
iodide,  the  change  would  be  expressed  by  the  following 
equation — 

Hg  +  I2  -  Hgl* 

This  means  that  one  atom  of  mercury  combines  with  two 
atoms  (one  molecule)  of  iodine,  and  forms  one  molecule  of 
mercuric  iodide,  which  contains  one  atom  of  mercury  and  two 
atoms  of  iodine. 

By  means  of  similar  equations  we  can  express  all  chemical 
changes  which  are  understood.  Take,  for  example,  the  various 
reactions  by  which  the  element  hydrogen  was  obtained, 
Chapter  VI. 

The  action  of  sodium  on  water  is  represented  thus — 

H2O  +  Na  =  NaHO  +  H 

which  means  that  when  one  molecule  of  water  is  acted  on  by 
one  atom  of  sodium,  one  atom  of  hydrogen  is  displaced,  and 
a  molecule  of  sodium  hydroxide  is  formed.  Or,  again,  the 
equation 

H2S04  +  Zn  =  ZnS04  +  H2 

means  that  one  atom  of  zinc  acts  on  one  molecule  of  sulphuric 
acid,  displacing  two  atoms  (one  molecule)  of  hydrogen  and 
forming  a  molecule  of  zinc  sulphate. 


Chemical  Equations.  155 

As  matter  can  neither  be  created  nor  annihilated,  every 
single  atom  which  figures  on  the  left  side  of  an  equation  must 
be  accounted  for  on  the  other  side.  In  no  chemical  change  is 
there  such  a  thing  as  an  atom  becoming  lost  or  destroyed,  they 
are  only  made  to  enter  into  fresh  combinations.  Before  we 
can  correctly  write  down  any  equation,  therefore,  we  must 
know  what  these  new  combinations  are.  For  example,  when 
sodium  chloride,  sulphuric  acid,  and  manganese  dioxide  are 
mixed  together,  a  chemical  change  takes  place, 

2NaCl  +  2H2SO4  -f  MnO,  = 

If  we  do  not  know  what  compounds  are  actually  produced, 
we  cannot  complete  the  equation.  We  must  not  simply  re- 
arrange the  symbols,  even  although  we  only  make  use  of  the 
same  number  of  the  identical  symbols  which  are  here  given. 
There  might  be  fifty  possible  ways  of  arranging  the  symbols, 
but  there  is  only  one  which  represents  what  actually  happens 
when  chemical  action  takes  place  between  these  compounds. 

The  three  Modes  of  Chemical  Action.— It  was 
stated  on  p.  14  that  chemical  action  takes  place  according 
to  three  general  modes ;  these  are — 

(i)  By  the  direct  union  of  two  molecules  to  form  a  more 
complex  molecule. 

Experiment  129. — Fill  a  glass  cylinder  with  gaseous  hydro- 
chloric acid  (as  described  in  Exp.  104),  and  cover  it  with  a  glass 
plate.  Take  another  similar  cylinder  and  pour  into  it  a  few  drops 
of  the  strongest  ammonia  ;  allow  the  liquid  to  run  round  the  glass, 
and  pour  out  any  excess.  This  vessel  will  now  contain  a  quantity 
of  ammonia  gas.  This  might  be  proved  by  carefully  smelling  the 
contents  of  the  cylinder. 

Now  invert  one  cylinder  upon  the  other,  mouth  to  mouth,  and 
notice  that  the  two  colourless  gases  combine  to  form  a  white  solid, 
which  appears  as  a  cloud  or  smoke. 

This  solid  is  ammonium  chloride  ;  the  molecules  of  the  two 
gases,  hydrochloric  acid  and  ammonia,  have  united  to  form  the 
more  complex  molecules  of  ammonium  chloride.  This  is  expressed 
by  the  equation — 

NHa  -i-  HC1  =  NH4C1. 


156     Chemical  Symbols,  Formula  and  Equations. 

(2)  By  the  exchange  of  atoms  in  different  molecules. 

Experiment  130. — Dissolve  a  small  quantity  of  mercuric  chloride 
in  water,  in  a  test-tube  (this  substance  dissolves  very  slowly  in  cold, 
but  more  readily  in  hot  water),  and  dip  into  the  solution  a  strip  of 
clean  copper  foil.  Notice  that  at  once  the  reddish  copper  becomes 
coated  with  a  film  of  mercury,  which  makes  it  look  as  though  it  had 
been  silvered.  The  mercury  in  the  solution  of  the  mercuric  chloride 
has  been  displaced  by  the  copper,  and  an  amount  of  copper  equiva- 
lent to  this  quantity  of  mercury  has  gone  into  the  solution.  We 
express  this  change  as  follows — 

HgCl2        +        Cu  CuCl2        +        Hg. 

Mercuric  chloride.  Copper  chloride. 

If  the  action  is  allowed  to  go  on  long  enough,  the  whole  of 
the  mercury  will  be  thrown  out  or  displaced  by  copper. 

We  can  easily  convince  ourselves  that  the  solution  contains 
copper  chloride  by  making  three  little  tests. 

Experiment  131. — (a)  Pour  a  few  drops  of  copper  chloride 
solution  into  water  in  a  test-tube,  and  add  ammonia.  Notice  the 
deep-blue  colour  produced.  This  is  a  characteristic  test  for  copper. 

(b}  Add  ammonia  to  a  dilute  solution  of  mercuric  chloride. 
Observe  that  a  white  precipitate  is  formed,  but  no  blue  colour. 

(c)  Now  add  ammonia  to  the  mercuric  chloride  solution  in 
which  the  copper  has  been  placed  for  a  short  time  and  then 
removed.  Notice  that  a  white  precipitate  is  produced,  and  also 
that  the  solution  turns  blue. 

Experiment  132. — Pour  a  few  drops  of  hydrochloric  acid  solu- 
tion into  water  in  a  test-tube,  and  add  a  little  silver  nitrate  solution. 
A  white  precipitate  at  once  forms.  The  silver  in  the  silver  nitrate 
has  changed  places  with  the  hydrogen  of  the  hydrochloric  acid 
according  to  the  equation — 

AgNO3      +      HC1      =  •    HNO?      +      AgCL 

Silver  nitrate.  Nitric  acid.         Silver  chloride. 

This  mutual  exchange  of  atoms,  which  takes  place  when 
both  the  reacting  substances  are  compounds,  is  often  spoken 
of  as  double  decomposition. 

(3)  By  the  rearrangement  of  atoms  within  the  molecule. — The 
elementary  student  will  not  meet  with  any  examples  of  chemical 
change  which  belong  to  this  class. 


Chemical  Equations. 


57 


As  a  further  example  of  the  use  of  chemical  equations, 
illustrations  may  be  given  of  the  formation  of  salts  by  the  action 
of  acids  upon  bases  (see  p.  69). 

(a)  With  metallic  oxides — 


ZnO 

Zinc  oxide. 


H2S04      = 


Sulphuric 
acid. 


H,O      4      ZnSO4 

Water.  Zinc  sulphate. 


PbO      4      2HNO3 

Lead  oxide.  Nitric  acid. 

(Litharge.) 


-      H,0 


Carbonic  acid. 


(b)  With  metallic  hydroxides — 
KHO       4       H2CO3       =       H2O 

Potassium 
hydroxide. 

CaH202 

Calcium  hydroxide 
(slaked  lime). 


2HC1      -      2H2O 

Hydrochloric 
acid. 


Pb(N03)2. 

Lead  nitrate. 


4       HKCO3 

Hydrogen  potassium 

carbonate. 
(Bicarbonate  of  potash.) 

4      CaCl2. 

Calcium  chloride 


(c)  With  ammonia — 

2NH3     4-      H2S04      -      (NH4)2S04 

Ammonia.  Sulphuric  acid.  Ammonium  sulphate. 


CHAPTER   XVII. 

THE   QUANTITATIVE   SIGNIFICANCE    OF   CHEMICAL    EQUATIONS. 

WE  have  seen  that  a  chemical  equation  tells  us  what  becomes 
of  the  various  substances  which  undergo  chemical  change.  By 
means  of  these  equations,  chemists  state  exactly  what  are  the 
various  substances  resulting  from  such  changes.  But  there  is 
still  more  information  conveyed  by  these  equations,  for  they 
actually  tell  us  the  exact  quantities  of  all  the  different  sub- 
stances which  undergo  change,  and  also  of  the  substances 
resulting  from  the  action. 

We  have  learnt  that  a  symbol  stands  for  an  atom,  and  also 
that  different  atoms  have  different  weights.  Thus,  the  symbol 
O  stands  for  one  atom  of  oxygen  ;  but  an  atom  of  oxygen  is 
sixteen  times  as  heavy  as  an  atom  of  hydrogen,  therefore  the 
symbol  O  stands  for  a  quantity  of  oxygen  which  is  sixteen 
times  as  heavy  as  the  quantity  of  hydrogen  represented  by  the 
symbol  H.  We  have  agreed  to  take  the  hydrogen  atom  as  the 
unit,  and  call  its  weight  i ;  H  therefore  stands  for  i  part  by 
weight  of  hydrogen,  and  O  for  16  parts  by  weight  of  hydrogen. 
In  the  same  way  the  symbols  Na,  Mg,  S,  stand  not  only  for 
sodium,  magnesium,  and  sulphur,  but  for  23  parts  by  weight 
of  sodium,  24  parts  by  weight  of  magnesium,  and  32  parts  by 
weight  of  sulphur.  This  fact  will  obviously  give  to  a  chemical 
equation  a  quantitative  meaning.  Thus,  the  equation— 

H2  +  O  =  H2O 

not  only  means  that  when  hydrogen  and  oxygen  combine, 
water  is  produced,  but  it  also  signifies  that  2  parts  by  weight 


Quantitative  Significance  of  Chemical  Equations     159 

of  hydrogen  unite  with  16  parts  by  weight  of  oxygen,  and  yield 
1 8  parts  by  weight  of  water — 

H2  +  0  =  H2O. 
2       16         18 

Or,  again,  the  equation — 

KC1O3  =  KC1  -f-  30, 

besides  meaning  that  when  potassium  chlorate  is  decomposed 
it  yields  potassium  chloride  and  oxygen,  has  also  the  quanti- 
tative significance — 

KC103  =  KC1  +  30 
f39  39         16  x  3 

35'5          35'5 
48 

122-5    =    74*5  +  48 

122-5  parts  of  potassium  chlorate  yield  74*5  parts  of  potassium 
chloride  and  48  parts  of  oxygen,  by  weight. 

Bearing  these  facts  in  mind,  it  is  quite  easy  to  calculate 
how  much  material  we  must  use  to  produce  any  required 
quantity  of  any  particular  product  of  a  chemical  change.  For 
instance,  suppose  we  wish  to  know  what  weight  of  oxygen 
could  be  produced  from,  say,  10  grams  of  potassium  chlorate. 
Then  from  the  above  equation  we  find  that  122*5  parts  by 
weight  of  potassium  chlorate  give  48  of  oxygen,  hence  we 
have  the  rule  of  three  sum — 

As  122-5   :    10    : :   48   :   x 

Grams  of  potassium  chlorate.      Grams  of  oxygen. 

48  X  10  , 

x  = —  =  3-9  grams  of  oxygen ; 

therefore  10  grams  of  potassium  chlorate  will  give  3*9  grams 
of  oxygen  gas. 

Again,  suppose  we  require  to  find  out  the  weight  of 
magnesium  oxide  which  is  produced  when  four  grams  of 
magnesium  are  burnt,  either  in  air  or  in  pure  oxygen.  We 


160     Quantitative  Significance  of  Chemical  Equations. 

first  write  the  equation  which  expresses  the  chemical  change 
that  takes  place — 

Mg  +  O  =  MgO 
24  +  16  =  40 

24  parts  by  weight  of  magnesium  give  40  parts  of  magnesium 
oxide,  how  much  will  4  grams  yield  ? 

:  •     40     :    x 

Grams  of  magnesium  oxide. 

x  -  4U  *  4  =  6*6  grams  of  magnesium  oxide. 
24 

Therefore,  4  grams  of  magnesium  on  burning,  will  form  6*6 
grams  of  magnesium  oxide. 

Let  us  take  one  more  example.  We  wish  to  make  10 
grams  of  hydrochloric  acid  gas,  What  weight  of  sodium  chloride 
and  of  sulphuric  acid  shall  we  require  ?  As  before,  we  write 
out  the  equation  for  the  preparation  of  hydrochloric  acid — 

NaCl  +  H2SO4  =  HNaSO4  +  HC1 
23  2  ii 

35'5        32  23  35-5 

64  32 

64 


58*5     4-    98        =         120        +        36-5 

36*5  parts  of  hydrochloric  acid  require  for  their  production 
58*5  parts  of  salt  and  98  parts  of  sulphuric  acid,  what  will  10 
grams  require  ? 

As  36  5    :    10    : :    58    :    x,  and 

Grams  of  hydrochloric  acid.  Grams  of  salt. 

as  36-5    :    10    : :    98    :    x' 

Grams  of  sulphuric  acid. 

Then  x  -  ^-    *  I0-  =  15*9  (nearly)  grams  of  salt. 
36'5 

And  x'  -  9    X  I0  =27-1  grams  of  sulphuric  acid. 
4      36*5 


Combination  of  Gases  by    Volume.  161 

Therefore,  in  order  to  prepare  10  grams  of  gaseous  hydrochloric 
acid,  we  require  15*9  grams  of  salt  and  27-1  grams  of  sulphuric 
acid 

The  method  by  which  we  are  able  to  calculate  the  volume 
which  any  given  weight  of  any  gas  will  occupy,  has  been  already 
explained  (p.  101)  and  exemplified  (p.  137),  and  since  we  are 
now  able  to  find  the  weight  of  gas  evolved  by  any  chemical 
process,  we  can  obviously  find  its  volume  also.  For  instance, 
suppose  we  require  to  know  the  volume  of  oxygen,  measured 
at  o°  and  760  mm.,  which  can  be  got  from  10  grams  of 
potassium  chlorate;  then  we  proceed  as  on  p.  159,  and  find 
first  the  weight  of  oxygen.  This  was  found  to  be  3-9  grams. 
Then  we  calculate  what  will  be  the  volume  at  N.T.P.  of  3*9 
grams  of  oxygen. 

Since  i  litre  of  oxygen  weighs  0*0896  x  16  =  i'433  grams, 
we  get  the  proportion — 

as  1^433    :   3-9    : :    i    :  x 

Grams  of  oxygen.          Litres  of  oxygen. 

Therefore  x  =  3'9  X  *  =27  litres  of  oxygen  at  N.T.P. 
r433 

If,  instead  of  at  N.T.P.,  we  wished  to  know  what  volume 
of  oxygen  measured  at  say  12°  C.  and  751  mm.  the  10  grams 
of  potassium  chlorate  would  yield,  then,  after  finding  first  the 
weight,  then  the  volume  at  N.  T.P.,  we  must  make  the  correction 
for  temperature  and  pressure,  thus — 

2^(273,+ 12)  x  76o  =  2.8  Htres  at  I2.  c  and 

273  x  751 

Combination  of  Gases  by  Volume. — By  Exp.  103  it 
was  shown  that  when  hydrogen  and  chlorine  combine,  they 
do  so  in  equal  volumes;  and  also  that  when  they  united, 
there  was  no  change  in  the  total  volume,  for  the  volume  of 
hydrochloric  acid  produced,  occupied  the  same  space  as  did 
the  mixed  gases  before  combination.  We  say,  therefore,  that 
one  volume  of  .hydrogen  combines  with  one  volume  of  chlorine 
and  forms  two  volumes  of  hydrochloric  acid. 

M 


1 62    Quantitative  Significance  of  Chemical  Equations. 

Again,  on  p.  77,  an  experiment  is  described  which  shows 
that  when  hydrogen  and  oxygen  unite,  they  do  so  in  the 
proportion  two  volumes  of  hydrogen  and  one  volume  of  oxygen. 
When  these  three  volumes  of  mixed  gases  combine,  however, 
they  do  not  give  three  volumes  of  steam ,  but  only  two  volumes. 

The  relation  between  the  volumes  of  gases  which  combine, 
and  the  volumes  of  the  compounds  produced,  is  found  to  be 
equally  simple  in  all  cases,  and  is  expressed  in  the  general 
statement  known  as  the  Law  of  Gay-Lussac. 

When  chemical  action  takes  place  between  gases,  the  volume 
of  the  gaseous  product  bears  a  simple  relation  to  the  volumes  of 
the  reacting  gases. 

Avogadro's  hypothesis. — In  the  year  1811,  the  Italian 
physicist  Avogadro  advanced  a  theory  in  order  to  explain  the 
known  facts  concerning  the  behaviour  of  gases.  This  theory 
is  accepted  by  all  chemists  and  physicists,  and  is  called 
Avogadro's  hypothesis.  It  may  be  thus  stated  :  Equal  volumes 
of  all  gases  (under  the  same  conditions  as  to  temperature  and 
pressure}  contain  the  same  number  of  molecules. 

If  a  litre  of  oxygen  contains  the  same  number  of  molecules 
as  a  litre  of  hydrogen,  it  will  be  evident  that,  if  we  weigh  a 
litre  of  each  gas,  the  difference  between  these  weights  will  also 
stand  for  the  difference  between  the  weights  of  a  single  molecule 
of  these  two  gases.  For  example,  let  us  suppose,  just  for  the 
sake  of  argument,  that  a  litre  of  oxygen  contains  a  million 
oxygen  molecules,  and  that  on  weighing  this  volume  we  find  it 
weighs  1*4336  grams.  According  to  Avogadro,  a  litre  of 
hydrogen  would  also  contain  a  million  molecules ;  suppose  we 
find  that  this  volume  of  hydrogen  weighs  0*0896  grams.  These 
numbers  then  express  the  ratio  between  the  weight  of  a  million 
oxygen  molecules  and  a  million  hydrogen  molecules.  1*4336 
is  sixteen  times  as  much  as  0-0896, 

for  1*4336  :  0*0896  : :  16  :  i 

If,  therefore,  a  million  oxygen  molecules  weigh  sixteen  times 
as  much  as  a  million  hydrogen  molecules,  obviously  each 
oxygen  molecule  must  be  sixteen  times  as  heavy  as  each 
molecule  of  hydrogen. 


Molecular   Weights  of  Gases.  163 

Therefore,  in  order  to  find  out  how  much  heavier  any 
gaseous  molecule  is  than  a  molecule  of  hydrogen,  all  that  is 
necessary  is  to  compare  the  weights  of  equal  volumes  of  the 
particular  gas  and  hydrogen. 

The  Densities  of  Gases. — The  density  of  a  gas  is  the 
weight  of  it  as  compared  with  the  weight  of  an  equal  volume 
of  some  other  gas  which  is  chosen  as  the  standard.  Some- 
times air  is  taken  as  the  standard.  Then,  when  we  say  that  the 
density  of  chlorine,  for  example,  is  2*45,  we  mean  that  a  given 
volume  of  this  gas  is  2*45  times  as  heavy  as  an  equal  volume 
of  air.  More  generally  hydrogen  is  the  standard,  in  which 
case  the  density  of  chlorine  is  35*5;  that  is  to  say,  a  given 
volume  of  the  gas  is  35-5  times  as  heavy  as  an  equal  volume 
of  hydrogen. 

Air  is  14*44  times  as  heavy  as  hydrogen;  therefore,  if  we 
know  the  density  of  any  gas  with  reference  to  one  standard,  we 
can  readily  calculate  by  "  rule  of  three *'  what  is  its  density  in 
terms  of  the  other;  thus,  the  density  of  chlorine  is  35-5  com- 
pared with  hydrogen,  what  is  its  density,  air  =  i  ? 

14-44  :  i  ::  35-5  \x 
^35-SXi 
14-44 

When  we  have  ascertained  the  density  of  a  gas  as  compared 
with  hydrogen,  we  have  also  found  (according  to  Avogadro's 
hypothesis)  the  ratio  between  the  weight  of  a  molecule  of  that 
gas  and  a  molecule  of  hydrogen. 

Molecular  Weights  of  Gases. — The  weight  of  a  mole- 
cule of  any  gas,  as  compared  with  the  weight  of  one  atom  of 
hydrogen,  is  called  the  molecular  weight  of  that  gas,  and  as 
the  molecule  of  hydrogen  consists  of  two  atoms,  therefore 
the  molecular  weight  of  a  gas  must  be  double  its  density. 

For  instance,  the  density  of  chlorine  is  found  by  experi- 
ment to  be  35 '5  ;  that  is,  its  molecule  is  35*5  times  as  heavy  as 
a  molecule  of  hydrogen ;  it  will,  therefore,  obviously  be  seventy- 
one  times  as  heavy  as  half  a  molecule  of  hydrogen  (or  one 
atom  of  hydrogen).  The  unit  or  standard  of  comparison  for 


164    Quantitative  Significance  of  Chemical  Equations. 

densities    is    the    molecule   of   hydrogen,  while    the    unit    for 
molecular  weights  is  the  half  molecule,  or  atom  of  hydrogen. 

By  experimentally  finding  the  densities  of  gases,  and,  there- 
fore, learning  their  molecular  weights,  chemists  are  able  to 
ascertain  the  atomic  weights  of  some  of  the  elements ;  for  the 
smallest  weight  of  an  element  that  is  ever  found  in  a  volume 
of  gas  equal  to  the  volume  occupied  by  one  molecule  of 
hydrogen,  is  taken  as  the  atomic  weight  of  that  element. 

The  Unit  Volume. — If  Avogadro's  hypothesis  be  true,  if 
equal  volumes  of  all  gases  contain  (under  the  same  conditions 
of  temperature  and  pressure)  the  same  number  of  molecules,  it 
follows  that  the  molecules  of  all  gases  occupy  the  same 
volume ;  or,  in  other  words,  they  occupy  the  same  volume  as 
a  molecule  of  hydrogen.  Each  molecule  of  oxygen,  or 
chlorine,  or  hydrochloric  acid  gas,  or  water  gas  (steam), 
occupies  the  same  space  as  a  hydrogen  molecule. 

The  space  occupied  by  a  molecule  of  hydrogen  is  called 
two  unit  volumes  (the  unit  volume  being  taken  as  the  space 
occupied  by  an  atom  of  hydrogen,  or  the  atomic  volume  of 
hydrogen),  hence  we  can  say  that  a  molecule  of  all  gases 
occupies  two  unit  volumes. 

If  we  bear  in  mind  that  gaseous  molecules  have  the  same 
volume,  and  if  we  write  our  chemical  equations  so  as  to 
represent  complete  molecules,  then  the  equation  at  once  tells 
us  the  volumetric  relations  between  the  reacting  gases  and  the 
gaseous  products  of  the  change.  For  instance,  the  equation — 

H2  +  Cla  =  2HC1 

not  only  carries  the  information  that  hydrogen  and  chlorine 
combine  to  form  hydrochloric  acid ;  and  that  the  proportion 
by  weight  is  i  part  of  hydrogen  to  35*5  parts  of  chlorine, 
yielding  36*5  parts  of  hydrochloric  acid;  but  it  tells  us  that 
i  molecule  (2  unit  volumes)  of  hydrogen,  combine  with  i 
molecule  (2  unit  volumes)  of  chlorine,  and  give  2  molecules 
(4  unit  volumes)  of  hydrochloric  acid  gas.  We  see  at  once, 
therefore,  from  the  equation  that  when  these  elements  combine 
there  is  no  change  of  volume ;  4  unit  volumes  of  the  mixed 
gases  giving  4  unit  volumes  of  the  compound. 


The  Unit   Volume.  165 

Again,  in  the  equation — 

2H2  -f  (X  =  2H20 

we  see  that  2  molecules  (4  unit  volumes)  of  hydrogen  combine 
with  i  molecule  (2  unit  volumes)  of  oxygen,  and  give  2 
molecules  of  water,  which,  if  measured  in  the  gaseous  state, 
occupy  4  unit  volumes.  (These  volume  relations  only  apply 
to  substances  in  the  gaseous  condition.)  At  a  glance,  therefore, 
we  see  from  this  equation  that  when  hydrogen  and  oxygen 
unite  to  form  steam,  there  is  a  contraction  in  the  volume, 
equal  to  one-third  of  the  original. 


CHAPTER  XVIII. 

THE   ATMOSPHERE. 

IN  the  early  days  of  Chemistry,  the  word  air  was  used  for  all 
gases.  Hydrogen  was  called  inflammable  air  ;  oxygen,  dephlogis- 
tigated  air,  and  so  on.  At  the  present  time  we  employ  the 
word  only  to  denote  the  gas  which  surrounds  and  envelopes 
the  earth. 

Formerly,  air  was  regarded  as  one  of  the  four  so-called 
elements — earth,  air,  fire,  and  water  ;  now  we  know  that  air  is 
in  no  sense  an  element,  neither  is  it  a  compound,  but  a 
mixture  of  several  gases,  some  of  which  are  elements  and  some 
compounds. 

The  exact  composition  of  the  atmosphere  has  been  made 
the  subject  of  close  investigation  by  many  chemists,  notably 
Lavoisier,  Boyle,  Priestley,  Cavendish,  and  Bunsen ;  so  that  it 
has  for  a  long  time  been  supposed  that  all  that  there  was  to 
know  about  the  composition  of  the  air  was  thoroughly  known. 
Nevertheless,  so  recently  as  1894  it  was  discovered  that  there 
is  a  certain  gas  present  in  the  air  which  had  hitherto  escaped 
observation  and  been  entirely  overlooked.  This  gas  has  been 
named  Argon. 

The  Two  chief  constituents  of  the  Air  are  Oxygen 
and  Nitrogen ;  all  the  others  are  present  in  comparatively 
minute  quantities.  By  various  experiments  we  have  learnt 
that  when  substances  burn  in  the  air,  they  are  in  reality 
combining  with  the  oxygen  present ;  we  can,  therefore,  make 
use  of  this  fact  in  order  to  remove  this  constituent  from  the 
atmosphere. 


The  Atmosphere.  167 

Experiment  133.— Place  a  small  bit  of  phosphorus,  which  has 
been  wiped  dry  with  blotting-paper,  in  a  little  porcelain  dish 
floating  on  water  in  a  pneumatic  trough.  Light  the  phosphorus 
and  stand  a  wide-mouthed  bottle  or  cylinder  over  it.  (See  Exp.  70, 
Fig.  51.)  As  the  phosphorus  continues  burning  in  the  enclosed 
space,  the  flame  gradually  becomes  fainter,  until  finally  it  goes  out. 
Now  notice  that  the  water  rises  inside  the  vessel,  showing  that  a 
part  of  the  air  which  originally  filled  it  has  disappeared.  Allow  the 
jar  to  remain  a  little  while,  so  that  the  white  fumes  of  phosphorus 
pentoxide  may  have  time  to  clear  away,  by  getting  dissolved  in  the 
water,  and  it  will  be  seen  that  the  remaining  gas  is  quite  clear 
and  colourless.  Of  course,  if  phosphorus  would  not  go  on  burning 
in  this  gas  which  is  left,  it  is  hardly  to  be  expected  that  a  taper 
would  burn  in  it ;  but,  to  make  sure,  the  jar  may  be  removed  (by 
slipping  a  glass  plate  beneath  its  mouth  and  lifting  it  out  of  the 
trough)  and  a  lighted  taper  or  candle  lowered  into  the  gas. 

If  this  experiment  be  made  in  a  different  way,  we  can 
roughly  measure  what  volume  of  oxygen  there  is  in  air. 

Experiment  134. — Take  a  long  wide  glass  tube,  sealed  up  at  one 
end,  and  fitted  with  a  cork  at  the  other.  Divide  the  tube  into  five 
equal  divisions,  which  can  be  marked  by  slipping  indiarubber  rings 
over  the  tube.  Drop  a  dry  piece  of  phosphorus  into  the  tube  and 
cork  it  up.  Next  dip  the  end  into  warm  water,  so  as  to  melt  the 
phosphorus,  and  make  it  begin  to  burn,  and  then  quickly  tip 
the  tube  up  so  as  to  allow  the  burning  phosphorus  to  run  down  the 
inside.  In  this  way  it  immediately  combines  with  all  the  oxygen 
present.  After  a  moment,  when  the  tube  has  cooled,  open  the  end 
under  water  in  a  trough.  Notice  that  the  water  rises  to  the  first 
ring,  showing  that  \  of  the  air  has  been  taken  out  by  the 
phosphorus. 

Therefore,  -J-  of  the  air  consists  of  oxygen,  which  is  one  of 
the  two  chief  constituents.  The  remaining  £  consists  of  the 
other  chief  ingredient,  namely,  nitrogen^  mixed  with  the  other 
gases  which  are  present  in  very  small  quantities. 

The  average  of  a  large  number  of  experiments,  shows  that 
the  proportion  of  oxygen  and  nitrogen  in  the  air  is  very  nearly 
21  parts  of  oxygen  to  79  of  nitrogen,  by  volume. 

The  gases  which  are  present  in  small  quantities  are  water 
vapour  (to  a  variable  extent),  argon  (about  0*8  per  cent,) 


1 68  The  Atmosphere. 

carbon   dioxide    (about   0*04    per   cent.),    and    still    smaller 
quantities  of  ozone,  ammonia,  and  nitric  acid. 

The  Composition  of  Air  by  Weight  was  carefully 
determined  by  Dumas  in  1841.  In  order  to  understand  his 
method,  we  may  make  the  following  experiment :  — 

Experiment  135. — Heat  a  quantity  of  bright  copper  gauze  in  a 
piece  of  combustion  tube,  in  a  gas  furnace,  and  pass  a  slow  stream 
of  air  from  a  gas  holder  through  the  tube.  Notice  that  the  copper 
becomes  black,  and  begins  to  do  so  first  at  the  end  at  which  the  air 
enters.  Collect  the  gas  which  passes  out  of  the  tube,  in  the 
pneumatic  trough,  and  test  it  by  dipping  a  lighted  taper  into  a  jar 
of  it.  Note  that  it  behaves  exactly  as  the  nitrogen  obtained  from 
air  in  Exp.  133.  The  copper  in  this  experiment  has  combined  with 
the  oxygen  of  the  air,  leaving  the  nitrogen. 

Dumas  heated  a  weighed  quantity  of  copper  in  a  tube,  one 
end  of  which  was  connected  to  a  large  vacuous  globe  with  a 
stopcock,  previously  carefully  weighed.  On  slightly  opening 
the  tap,  air  was  slowly  drawn  over  the  heated  copper.  The 
copper  combined  with  the  oxygen,  and  the  nitrogen  was 
received  in  the  globe.  In  order  to  remove  the  other  gases 
which  were  present  in  small  quantities,  the  air  was  first  made 
to  pass  through  a  series  of  U  tubes,  filled  with  materials  which 
would  absorb  the  water  vapour  and  the  carbon  dioxide.  Argon 
was  unknown,  and,  therefore,  went  in  with  the  nitrogen.  The 
average  of  a  number  of  experiments,  showed  that  the  proportion 
of  oxygen  and  nitrogen  by  weight  was,  oxygen  23  parts,  and 
nitrogen  77  parts. 

Air  a  Mixture. — The  oxygen  and  nitrogen  in  the  air  are 
not  chemically  combined,  but  only  mixed  together.  The  chief 
proofs  of  this  are  the  following  : — 

(1)  The  composition  of  the  air  is  not  constant,  the  propor- 
tions of  oxygen  and  nitrogen  are  found  to  vary  slightly.     If 
the  air  was  a  chemical  compound  of  these  gases,  then,  accord- 
ing to  the  law  of  constant  composition,  the  constituents  would 
always  be  present  in  exactly  the  same  proportions. 

(2)  The  proportion  of  oxygen  and  nitrogen  in  the  air  does 
not  bear  any  simple  relation  to  the  atomic  weights  of  these 
elements. 


Air  a  Mixture.  169 

(3)  The  two  gases  can  be  separated  by  mechanical  means ; 
such  as  by  dissolving  in  water  (oxygen  is  more  soluble  than 
nitrogen),  and  by  diffusion. 

Diffusion  of  Gases. — The  gases  in  the  air  have  different 
densities;  water  vapour  =  8,  nitrogen  =  14,  oxygen  =  16, 
carbon  dioxide  =  22.  Why  is  it  that  the  heavy  carbon  dioxide 
does  not  sink  to  the  ground  and  form  a  bottom  layer,  and  the 
light  water  vapour  rise  up  above  the  oxygen  and  nitrogen? 
Two  chief  causes  operate  to  prevent  this,  and  to  keep  the 
gases  thoroughly  mixed.  The  first  is  wind  and  air  currents  • 
the  second  is  the  property,  belonging  to  all  gases,  of  diffusion. 

If  we  take  a  jar  filled  with  hydrogen  and  hold  it  mouth 
downwards,  although  the  gas  is  fourteen  times  lighter  than  air, 
it  will  nevertheless  make  its  escape  downwards  out  of  the 
bottle. 

Experiment  136. — Take  two  soda-water  bottles,  and  fill  one  with 
hydrogen  and  the  other  with  oxygen.  Join  the  bottles  by  means 
of  a  long  piece  of  combustion  tube  passing  through  two  corks  which 
fit  the  bottles.  Stand  the  apparatus  upright,  the  oxygen  being  at 
the  bottom,  and  leave  it  for  a  few  hours.  The  light  hydrogen  will 
gradually  find  its  way  down  the  tube  into  the  oxygen  bottle,  and 
the  oxygen  (sixteen  times  as  heavy)  will  pass  up  into  the  hydrogen 
bottle;  so  that,  after  a  time,  there  will  be  a  perfect 
mixture  of  the  two  gases  in  both  bottles.  That  the  /| 
gases  have  thus  mixed  can  be  proved  by  removing  the 
bottles  and  applying  a  light  to  their  mouths.  If  the  gas  \  I 
explodes,  it  shows  that  the  two  have  mixed.  If/ 

This  power  that  gases  have  of  mixing  them- 
selves together  is  called  diffusion.  All  gases  do  not 
move  or  diffuse  at  the  same  rate.  We  can  prove 
this  in  the  following  way. 

Experiment  137. — Take  a  short  clay  tobacco-pipe, 
and  cement  over  the   mouth  of  the   bowl  a  piece  of     F       8 
cardboard,  or  better,  a  thin  piece  of  unglazed  earthen- 
ware.    Attach  the  stem  of  the  pipe,  with  a  piece  of  indiarubber 
pipe,  to  a  U-shaped  glass  tube,  containing  some  coloured  water, 
as  shown  in   Fig.  78.      Now  fill  a  beaker  with  hydrogen,  hold 
its  mouth    downwards,   and   put   the    tobacco-pipe  up    into  the 


170  The  Atmosphere. 

beaker.  Notice  that  the  water  in  the  U-tube  is  immediately 
driven  down  the  limb  to  which  the  pipe  is  attached.  This  is 
because  the  light  gas  hydrogen  diffuses  through  the  porous  clay 
pipe  much  faster  than  the  air  inside  the  pipe  can  make  its  way  out, 
and  therefore  an  excess  of  gas  collects  inside,  and  pushes  down  the 
water.  Now  remove  the  beaker,  and  notice  that  the  water  not  only 
returns  to  the  level,  but  rises  considerably  in  the  limb  attached  to 
the  pipe.  The  condition  of  things  is  now  reversed,  the  hydrogen 
which  had  gone  inside  the  pipe  is  making  its  way  out  again  faster 
than  air  can  get  in,  hence  there  is  a  diminution  in  the  volume  of 
gas  inside. 

The  lighter  a  gas  is,  the  faster  does  it  diffuse. 

Graham's  Law. — Graham  found  out  the  law  which  regu- 
lates gaseous  diffusion,  which  may  be  thus  stated :  The  relative 
rates  of  diffusion  of  any  two  gases  are  inversely  as  the  square  roots 
of  their  densities.  For  example — 

The  density  of  hydrogen  is  i ;  the  square  root  of  i  =  i. 

The  density  of  air  is  14*44;  the  square  root  of  14*44  =  3'8- 

T,,       r       the  rate  of  diffusion    the  rate  of  diffusion  .       0  . 
i  nereiore          /•  i    j  •  r    •  •  •  3  <*  •  i 

of  hydrogen  of  air 

in  other  words,  hydrogen  diffuses  3*8  times  as  quickly  as 
air  does. 

Now  since  nitrogen  is  just  a  little  lighter  than  oxygen,  if 
we  pass  a  mixture  of  these  two  gases  through  a  long  porous 
pipe  (under  suitable  experimental  conditions)  nitrogen  will 
escape  through  the  walls  of  the  pipe  a  little  faster  than  oxygen, 
therefore  the  gas  which  is  delivered  out  at  the  end  of  the  pipes 
will  contain  rather  more  oxygen  in  proportion  to  nitrogen  than 
that  which  is  sent  in. 

By  driving  a  stream  of  air  through  such  pipes,  it  is  found 
in  like  manner,  that  what  comes  out  at  the  end  is  a  little  richer 
in  oxygen  than  the  air  which  is  sent  in.  Therefore  the  oxygen 
and  nitrogen  in  air  can  only  have  been  mixed  and  not  com- 
bined, or  it  would  not  be  possible  to  sift  one  away  from  the 
other  by  this  process  of  diffusion. 

Combustion. — We  know  that  a  great  many  substances 
will  burn  in  the  air.  It  is  customary  to  call  such  things  com- 
bustibles >  and  to  speak  of  the  air  as  a  supporter  of  combustion. 


Combustion.  17 1 

We  have  learnt  by  numerous  experiments  that  it  is  because 
of  the  oxygen  present,  that  the  air  supports  the  combustion  of 
burning  bodies ;  oxygen,  as  we  have  seen,  being  such  a  good 
supporter  of  combustion.  In  ordinary  language  we  call  a  gas 
a  supporter  of  combustion,  if  it  behaves  towards  common  com- 
bustibles in  the  same  way  that  air  does ;  but  in  reality  there  is 
no  distinction  between  a  combustible  and  a  supporter  of  com- 
bustion. For  instance,  we  have  seen  sulphur  burning  in 
oxygen,  and  we  therefore  call  sulphur  a  combustible,  and 
oxygen  the  supporter  of  combustion.  But  let  us  modify  the 
conditions  of  the  experiment. 

Experiment  138. — Heat  some  sulphur  in  a  wide  test-tube  until  it 
boils  and  the  dark  vapour  takes  fire  at  the  mouth.  Then  lower 
into  the  test-tube  a  bent  glass 
tube  (so  bent  that  it  can  enter 
the  test  tube)  through  which  a 
gentle  stream  of  oxygen  is  passing 
(Fig.  79).  As  the  jet  is  passed 
through  the  flame  of  sulphur 
burning  at  the  mouth,  the  oxygen 
is  ignited,  and  will  continue  burn- 
ing in  the  sulphur  vapour  when 
pushed  down  into  the  tube. 

Under  these  conditions  the 
oxygen  is  the  combustible,  and 
sulphur  vapour  is  the  supporter 
of  combustion.  Again,  we  have 
seen  hydrogen  burn  in  oxygen, 
but  we  can  reverse  the  condi- 
tions and  make  oxygen  burn  in 
hydrogen. 

Experiment  139. — Fill  a  jar 
with  hydrogen.  Hold  it  mouth 
downwards,  and  apply  a  light  to  ~~~FIG 

the  gas.  While  the  hydrogen  is 
burning,  thrust  a  jet,  from  which  oxygen  is  slowly  issuing,  up  into 
the  jar  (Fig.  80).  As  the  jet  passes  through  the  hydrogen  flame 
it  is  ignited,  and  goes  on  burning  just  as  it  did  in  the  sulphur 
vapour. 


172  The  Atmosphere. 

Here  oxygen  is  the  combustible,  and  hydrogen  the  supporter 
of  combustion. 

Combustion  is  merely  the  term  applied  to  describe  any 
chemical  action  which  takes  place  with  so  much  energy  as  to 
produce  light  and  heat.  All  the  more  common  cases  of  com- 
bustion, are  the  active  combination  of  substances  with  oxygen, 
that  is,  they  are  rapid  processes  of  oxidation ;  but  oxygen  is 
not  necessary  to  combustion ;  we  can  have  cases  of  combustion 
in  which  oxygen  does  not  participate,  for  we  know  of  many 
instances  in  which  chemical  action  takes  place  between  other 
substances  with  sufficient  energy  to  give  rise  to  light.  For 
instance — 

Experiment  140. — Take  a  jar  of  chlorine,  and  lower  into  it  a  jet 
from  which  ammonia  gas  is  slowly  escaping  (the  ammonia  being 
obtained  by  gently  warming  a  strong  solution 
of  ammonia,  in  a  little  flask;  see  Exp.  145). 
Notice  that  as  soon  as  the  jet  of  ammonia  enters 
the  chlorine  it  at  once  takes  fire,  without  being 
lighted,  and  goes  on  burning  in  the  chlorine, 
(Fig.  81).  There  is  here  no  oxygen  present ;  the 
ammonia  is  the  combustible,  and  the  chlorine  is 
the  supporter  of  combustion. 

We  have  also  seen  (Exp.  119)  that  many 
metals  take  fire  and  burn  when  brought  into 
chlorine. 

Temperature  of  Combustion.— The 
actual  temperature  which  is  produced  by  any 
particular  process  of  combustion  depends 
partly  upon  how  quickly  the  combustion 
proceeds.  For  instance,  if  we  burn  some 

FIG.  80.  .  .     -  .     -  . 

substance  m  oxygen  it  burns  much  faster  than 
when  burnt  in  air,  and  consequently  the  temperature  is  hotter. 

Experiment  141. — Burn  a  jet  of  hydrogen  from  an  oxyhydrogen 
blowpipe  (this  is  merely  a  very  fine-pointed  metal  tube  with  a  still 
smaller  one  passing  down  the  inside,  exactly  like  an  ordinary 
Herapath  blowpipe,  only  smaller).  Hold  a  little  piece  of  platinum 
wire  in  the  flame,  and  note  that  although  the  wire  becomes  very 
hot,  it  does  not  melt.  Next  hold  a  little  block  of  hard  lime  against 


Temperature  of  Combustion. 


173 


FIG.  81. 


the  flame  ;  the  lime  does  not  seem  to  get  very  hot.  Now  gently 
turn  on  the  oxygen,  so  that  the  hydrogen  flame  is  fed  with 
oxygen.  Notice  that  the 
flame  does  not  show  any 
more  light,  but  if  the  plati- 
num wire  is  held  in  the  flame 
it  is  instantly  melted,  be- 
cause the  flame  is  now  so 
much  hotter  ;  and  if  the 
lime  is  brought  into  the  flame 
it  at  once  gets  so  hot  as  to 
emit  a  very  bright  light. 
This  is  the  oxyhydrogen 
lime-light. 

In  everyday  life  when  ^ 
we  wish  to  increase  the  rate 
of  combustion,  and  conse- 
quently raise  the  temperature  of  combustion,  we  increase  the 
draught  of  air  (by  the  use  of  bellows,  for  instance)  so  that  more 
oxygen  is  driven  against  the  burning  body  in  a  given  time. 

Ignition  point. — The  particular  temperature  at  which  a 
substance  begins  to  burn, 
or  to  "  take  fire,"  is  called 
its  ignition  point.  Some- 
times this  is  lower  than  the 
ordinary  temperature  of 
the  room,  in  which  case  the 
substance  takes  fire  by  it- 
self when  brought  into  the 
air.  Obviously  such  things 
as  these  must  be  kept  so 
that  they  do  not  come  into 
contact  with  the  air. 

Experiment  142.— Place  :f| 
a  small  quantity  of  caustic  ™ 
soda  solution  in  a  test-tube,  ^ 

and  put  in  it  a  piece  of  phos- 
phorus about  the  size  of  a  pea.     Attach  a  cork  with  two  tubes, 
arranged  as  in  Fig.  82.     First  pass  a  stream  of  ordinary  coal  gas 


1/4  The  Atmosphere. 

through  the  little  apparatus  by  means  of  the  indiarubber  pipe  con- 
nected with  the  tube  T.  This  is  in  order  to  sweep  out  the  air  from 
the  apparatus.  Now  gently  boil  the  liquid,  and  a  gas  is  given  off 
called  phosphoretted  hydrogen,  which  will  bubble  through  the  water 
in  the  little  basin. 

In  a  minute  or  two,  when  the  coal  gas  has  been  expelled,  and 
only  the  phosphoretted  hydrogen  is  bubbling  out,  each  bubble  will 
take  fire  as  it  conies  into  the  air.  This  gas  has  a  very  low  igniting 
point. 

In  all  the  familiar  processes  of  combustion,  it  is  neces- 
sary to  first  heat  the  combustible  substance  in  order  to  start 
the  combustion ;  that  is  to  say,  the  igniting  point  is  above  the 
common  temperature,  and  therefore  the  substance  must  be 
heated  up  to  that  point  before  active  chemical  combination 
can  take  place.  If  after  a  substance  is  ignited,  it  will  continue 
burning  by  itself,  like  a  candle,  or  piece  of  paper,  this  shows 
that  the  temperature  of  combustion  is  higher  than  the  igniting 
point;  because,  as  each  particle  burns,  the  heat  given  out  is 
able  to  set  fire  to  the  next  particle,  and  so  on. 

Heat  of  combustion  is  the  amount  of  heat  as  dis- 
tinguished from  the  temperature^  produced  by  combustion.  If 
we  draw  a  pint  of  hot  water  and  a  gallon  of  water  from  the 
same  boiler,  the  temperature  of  each  sample  will  be  the  same ; 
a  thermometer  placed  in  each  will  show  the  same  temperature. 
But  it  is  obvious  that  the  amount  of  heat  in  the  gallon  of 
water  is  greater  than  in  the  pint;  there  is,  of  course,  eight 
times  as  much  heat  in  the  one  as  in  the  other.  From  this 
illustration  it  will  be  evident  that  the  amount  of  heat  cannot 
be  ascertained  by  the  thermometer.  It  is  really  measured  by 
finding  out  how  much  water  it  is  capable  of  heating  from 
o°  to  i°. 

That  amount  of  heat  which  will  raise  the  temperature  of 
i  gram  of  water  from  o°  to  i°  is  taken  as  the  unit,  and  is 
called  the  thermal  unit  (or  calorie). 

Now  the  amount  of  heat  produced  in  any  process  of  com- 
bustion, is  exactly  the  same  whether  the  process  be  slow  or 
quick,  therefore  it  is  quite  independent  of  the  temperature  that 
is  produced.  For  example,  the  amount  of  heat  given  out 


Heat  of  Combustion.  175 

when  a  definite  quantity  of  hydrogen  is  burnt  in  the  air,  is 
exactly  the  same  as  that  produced  when  the  same  quantity  of 
hydrogen  is  burnt  in  oxygen,  although,  as  we  have  seen,  the 
temperature  in  the  latter  case  is  very  much  higher.  Indeed, 
this  is  also  true  if  the  process  of  oxidation  is  so  slow  that 
there  is  no  active  combustion.  Thus,  when  a  quantity  of  iron 
slowly  rusts,  heat  is  produced ;  but  the  process  is  spread  over 
such  a  long  time,  that  the  iron  never  even  gets  warm,  the 
heat  being  conducted  away  as  fast  as  it  is  produced.  When 
the  same  quantity  of  iron  is  burnt  in  oxygen,  the  temperature 
rises  enormously,  because  the  process  is  complete  in  a  few 
moments,  but  the  actual  amount  of  heat  is  the  same  in  both 
cases. 


CHAPTER   XIX. 


NITROGEN    AND    ITS    COMPOUNDS. 

Nitrogen. — The  element  nitrogen,  as  we  have  seen,  is  present 
in  the  free  state  in  the  air ;  f  of  which  is  nitrogen.  It  can  be 
obtained  from  the  air  by  removing  the  oxygen  either  with  phos- 
phorus (Exp.  133)  or  by  means  of  heated  copper  (Exp.  135). 

Nitrogen,  combined  with  other  elements,  is  present  in  a 
number  of  compounds,  from  some  of  which  the  element  is 
readily  expelled. 

One  of  the  commonest  of  the  compounds  of  nitrogen  is 
ammonia ;  a  compound  of  nitrogen  with  hydrogen.  If  we  act 
on  this  compound  with  chlorine,  the  chlorine  takes  the 
hydrogen  (forming  hydrochloric  acid)  and  the  nitrogen  is  set 
free,  according  to  the  equation, 

2NH3  +  3C1.2  =  6HC1  +  N2 

Experiment  143. — Fit  a  wide-mouthed  bottle  with  a  cork  carry- 
ing two  wide  glass  tubes,  as  shown  in 
Fig.  83.  Half  fill  the  bottle  with  strong 
ammonia  solution,  and  pass  a  stream 
of  chlorine  (prepared  as  in  Exp.  no) 
through  the  liquid.  Notice  that,  as 
each  bubble  of  chlorine  enters  the  am- 
monia, there  is  a  flash  of  fire  in  the 
liquid,  so  energetic  is  the  chemical 
action  that  goes  on.  Observe  also  the 
white  fumes  ;  these  consist  of  ammo- 
nium chloride  ;  for,  as  soon  as  hydro- 
chloric is  formed  (as  shown  in  the 
equation),  it  combines  with  some  of  the 
ammonia  in  the  bottle  and  forms  ammonium  chloride.  Collect 
the  gas  which  passes  out,  in  the  pneumatic  trough.  The  object 


FIG.  83. 


Properties  of  Nitrogen.  177 

of  the  very  wide  tubes  is  that  they  may  not  get  stopped  up  by 
the  ammonium  chloride.  [Caution — do  not  let  the  experiment  go 
on  too  long  ;  as  soon  as  one  or  two  jars  of  gas  have  been  collected, 
stop  the  operation.] 

Another  compound  from  which  nitrogen  can  be  readily 
obtained  is  ammonium  nitrite,  N£T4NO2.  When  a  strong 
solution  of  this  salt  is  gently  heated,  it  splits  up  into  nitrogen 
and  water. 

NH4N02=N2-f-2H20. 

Experiment  144. — Put  about  equal  quantities  of  ammonium 
chloride  and  sodium  nitrite  into  a  flask,  fitted  with  a  cork  and 
delivery  tube,  and  about  one-third  fill  the  flask  with  water.  Gently 
heat  the  mixture  and  collect  the  gas  over  water.  As  soon  as  the 
action  begins,  remove  the  lamp,  and  allow  it  to  continue  by  itself. 
If  the  liquid  begins  to  boil  up  too  much,  cool  the  flask  by  bringing 
a  dish  of  cold  water  under  it. 

When  a  mixture  of  ammonium  chloride  and  sodium  nitrite 
is  heated,  the  two  salts  interact  on  each  other  and  give  sodium 
chloride  and  ammonium  nitrite,  thus — 

NH4C1  +  NaNO2  -  NaCl  +  NH4NO2, 

and  the  ammonium  nitrite  then  decomposes  according  to  the 
first  equation.  The  final  result,  therefore,  of  heating  these  two 
salts  together  is  expressed  thus — 

NH4C1  +  NaNO2  =  NaCl  +  2H2O  +  N2. 

Properties  of  Nitrogen. — Nitrogen  is  very  different  from 
any  of  the  gases  we  have  a*s  yet  studied.  When  a  taper  is  put 
into  it  we  see  that  the  gas  is  not  like  hydrogen,  for  it  will  not 
burn.  It  is  not  like  oxygen,  for  it  will  not  allow  the  taper  to 
burn,  but  at  once  extinguishes  it.  It  is  not  like  chlorine  or 
hydrochloric  acid.  The  only  property  of  this  gas  that  can 
easily  be  shown  is,  that  it  seems  to  have  no  properties.  It 
does  not  burn.  It  does  not  support  combustion.  It  is  not 
acid ;  does  not  bleach ;  does  not  act  on  metals ;  is  not 
poisonous.  Indeed  nitrogen  is  one  of  the  most  inactive  or 
inert  substances  we  know.  It  will  not  support  animal  life, 
not  because  it  is  in  any  way  injurious,  but  simply  because 


178  Nitrogen  and  its  Compounds. 

animals  must  have  free  oxygen  to  breathe ;  an  animal  placed 
in  nitrogen  dies  from  suffocation,  just  as  it  would  if  immersed 
in  water. 

Nitrogen  is  slowly  absorbed  by  red  hot  magnesium,  forming 
a  compound  of  nitrogen  and  magnesium.  This  is  one  method 
by  which  nitrogen  is  separated  from  argon,  which  is  even  more 
inert  a  substance  than  nitrogen  itself. 

Although  nitrogen  will  not  burn,  it  will  combine  with 
oxygen  slowly,  when  electric  sparks  are  passed  through  a 
mixture  of  the  two  gases.  In  this  respect  again  it  differs  from 
the  still  more  inactive  gas  argon,  which  does  not  unite  with 
oxygen  when  sparked  with  that  gas. 

EPITOME. 

Nitrogen  occurs  uncombined  in  the  air,  to  the  extent  of  about 
four-fifths.  It  is  obtained  from  air  by  withdrawing  the  oxygen, 
either  by  burning  phosphorus  or  red  hot  copper.  These  combine 
with  the  oxygen  and  leave  the  nitrogen. 

The  chemical  compound  from  which  nitrogen  is  prepared,  is 
ammonium  nitrite.  This  when  heated  gives  only  water  and 
nitrogen. 

Nitrogen  is  characterised  by  great  inertness.  It  is  a  colourless, 
odourless,  tasteless  gas  ;  does  not  burn,  nor  support  combustion  or 
respiration  ;  is  not  poisonous. 

It  does  not  easily  unite  directly  with  other  elements.  At  a  high 
temperature  it  combines  with  a  few  metals,  and  also  with  oxygen. 

Just  as  chlorine  is  a  member  of  a  certain  little  family  of  elements 
(see  The  Halogens,  p.  133),  so  nitrogen  is  also  the  representative  of 
another  group  or  family,  which  consists  of  the  five  elements 
nitrogen,  phosphorus,  arsenic,  antimony,  and  bismuth.  These  have 
very  little  likeness  to  each  other  in  their  outward  appearance,  but 
they  are  closely  related  in  their  chemical  behaviour. 

Nitrogen  and  phosphorus  are  true  non-metals,  one  being  a  gas, 
the  other  a  wax-like  solid  ;  they  have,  therefore,  no  properties  which 
belong  to  metals,  no  metallic  lustre,  no  power  of  conducting  heat 
or  electricity  ;  and  they  yield  oxides  which  are  acid-forming 

Antimony  and  bismuth  are  metals  ;  they  have  metallic  lustre, 
conduct  heat  and  electricity,  and  form  oxides  which  are  basic. 

Arsenic  stands  on  the  border  line  between  the  metals  and  non- 
metals,  and  is  called  a  metalloid.  It  is  a  black  shiny  solid,  with 


Compounds  of  Nitrogen.  179 

about  as  much  lustre  as  graphite.     It  conducts  heat  and  electricity, 
but  its  oxides  are  acid-forming  compounds. 

The  detailed  study  of  phosphorus,  arsenic,  antimony,  and 
bismuth  does  not  come  within  the  scope  of  this  elementary  book. 

Compounds  of  Nitrogen. — Although  nitrogen  in  the 
free  state  is  such  an  inert  element,  it  has  very  strong  chemical 
affinities  when  in  combination  with  other  elements.  It  forms 
a  number  of  compounds  with  oxygen,  with  hydrogen,  and 
with  both  oxygen  and  hydrogen  together.  It  is  also  one 
of  the  constituents  of  a  large  number  of  animal  and  vegetable 
substances,  where  it  is  associated  with  carbon,  hydrogen,  and 
oxygen. 

When  such  animal  substances  decay  or  "go  bad,"  one 
of  the  first  products  of  the  decomposition  is  ammonia.  The 
nitrogen  in  the  compound  combines  with  some  of  the  hydrogen 
and  forms  this  compound.  Hence  there  is  generally  a  strong 
smell  of  ammonia  in  stables  and  in  urinals,  where  nitrogenous 
animal  matter  is  undergoing  decomposition.  This  natural 
process  of  decomposition  is  imitated  artificially,  when  we  heat 
such  a  compound  so  as  not  to  allow  air  to  get  to  it,  so  that  it 
does  not  take  fire.  For  example,  when  coal  is  heated  in 
retorts,  as  in  the  manufacture  of  ordinary  coal-gas,  air  does 
not  get  to  the  coal,  and  therefore  it  does  not  burn  as  it  would 
in  an  open  fireplace,  but  it  is  decomposed  into  a  great  variety 
of  compounds,  some  solid,  some  liquid,  and  some  gaseous. 
The  process  of  heating  substances  in  this  manner  is  called 
destructive  distillation,  to  distinguish  it  from  the  ordinary 
operation  of  distilling  where  the  original  substance  is  not 
decomposed.  Now,  although  coal  is  very  largely  composed 
of  carbon,  it  also  contains  amongst  other  constituents  some 
nitrogen,  and  some  hydrogen,  and  therefore,  when  it  is 
destructively  distilled,  one  of  the  gaseous  products  formed  is 
ammonia.  This  dissolves  in  the  water,  which  is  another  product 
of  the  decomposition,  yielding  the  so-called  "ammoniacal 
liquor  "  of  the  gas-works.  This  is  the  chief  source  from  which 
ammonia  is  now  obtained. 

Another  substance  which  is  formed  when  animal  matter 
•containing  nitrogen  is  allowed  to  decompose  slowly  by  itself 


180  Nitrogen  and  its  Compounds. 

in  the  presence  of  air,  is  nitric  acid.  Sometimes  in  the  neigh- 
bourhood of  ill-drained  stables  or  dwellings,  especially  in  hot 
climates,  where  decomposing  animal  matter  soaks  into  the 
earth,  crystals  are  to  be  seen  on  the  soil  or  lower  parts  of  the 
walls.  These  crystals  consist  of  nitre.  The  nitric  acid  formed 
by  the  decomposition  of  the  organic  matter,  combines  with 
potash  present  in  the  soil  and  forms  potassium  nitrate  (saltpetre 
or  nitre}. 

At  one  time  all  our  supplies  of  nitre  were  obtained  by  this 
process,  which  was  carried  on  by  purposely  mixing  manure 
and  such  decomposing  refuse  with  wood  ashes  (which  contain 
a  large  quantity  of  potash)  and  earth,  and  allowing  the  heaps 
to  remain  exposed  to  the  air,  occasionally  moistening  them 
with  drainage  from  manure. 

Nowadays  this  process  is  not  much  used,  because  enormous 
beds  of  sodium  nitrate  (called  Chili  saltpetre]  have  been  found, 
from  which  potassium  nitrate  can  easily  be  made. 

Ammonia  and  nitre  are  two  of  the  most  important  com- 
pounds of  nitrogen. 

Ammonia. — This  substance  is  a  gas,  but  long  before  it 
was  known  to  be  a  gas,  a  solution  of  it  in  water  was  known, 
and  was  called  spirits  of  hartshorn.  It  got  this  name  because 
it  was  obtained  by  the  destructive  distillation  of  horns  and 
hoofs  of  animals.  At  the  present  day  the  source  of  all  our 
ammonia  is  the  ammoniacal  liquor  of  the  gas-works.  This 
is  practically  a  solution  of  the  gas  in  tarry  water.  The  liquid 
used  in  the  laboratory,  and  called  ammonia,  is  simply  a 
solution  of  the  gas  in  pure  water.  If  we  heat  such  a  solution 
the  gas  is  all  expelled,  and  can  be  collected. 

Experiment  145. — Gently  heat  a  little  strong  solution  of  ammonia 
in  a  flask  provided  with  a  cork  and  short  exit  tube  bent  at  right 
angles,  and  attach  this  to  the  apparatus  for  collecting  gas  by 
upward  displacement  (Fig.  84).  Notice  that  the  ammonia  solution 
gives  off  its  gas  so  rapidly  that  it  appears  to  be  boiling,  although 
scarcely  warmed  to  the  temperature  of  the  hand.  In  a  few  minutes 
the  cylinder  will  be  rilled  with  ammonia  gas.  Now  remove  it  and 
place  it  mouth  downwards  in  a  trough  of  water.  Notice  that  the 
water  quickly  rises  in  the  cylinder,  showing  that  the  gas  is  far 


A  mmonia. 


181 


FIG.  84. 


too   soluble  in  water  to   allow  of  its  being  collected   over  that 
liquid. 

Experiment  146. — Collect  another  cylinder  of  the  gas,  and  as 
the  glass  is  filling,  hold  a  piece  of  litmus  paper  which  has  been 
reddened  (by  being  dipped  into  very 
dilute  acid)  against  the  exit  tube ;  ob- 
serve that  the  gas  is  strongly  alkaline. 
Also  let  the  gas  blow  against  a  moist- 
ened piece  of  turmeric  paper,  and  obtain 
the  reddish  brown  stain  due  to  the  action 
of  an  alkali  upon  it.  In  the  old  days 
when  all  gases  were  called  airs,  ammonia 
was  distinguished  as  the  alkaline  air. 
Cautiously  smell  the  gas,  not  by  applying 
the  nose  to  the  ttibe — this  would  give  too 
strong  a  sniff  of  it  and  would  be 
dangerous — but  by  gently  wafting  the 
escaping  gas  towards  the  face  with  the 
hand. 

Gently  thrust  a  lighted  taper  up  into 
the  gas.  Carefully  note  the  behaviour 
of  this  gas  towards  combustibles.  When  the  taper  is  plunged  right 
into  the  gas,  the  flame  is  extinguished,  and  the  ammonia  does  not 
take  fire.  But  when  the  taper  flame  is  cautiously  brought  into  the 
gas,  at  first  the  gas  seems  to  be  trying 'to  burn  ;  a  curious  brownish- 
yellow  flame  appears  to  surround  the  taper  flame  for  a  moment 
before  the  latter  is  put  out. 

Ammonia,  therefore,  will  not  bum  in  ordinary  air,  although 
it  seems  very  nearly  to  do  so.  But  if  we  add  a  little  oxygen 
to  the  air,  then  the  ammonia  will  bum  quite  easily. 

Experiment  147. — Close  one  end  of  a  wide  glass  tube  (a  gas- 
lamp  chimney)  with  a  cork  through  which  two  tubes  pass  :  a 
moderately  wide  one  reaching  to  the  top  of  the  chimney,  and  a 
narrow  one  passing  just  through  the  cork,  as  in  Fig.  85.  The 
wider  of  these  tubes  is  attached  to  a  small  flask  in  which  strong 
ammonia  solution  is  gently  heated  ;  the  other  is  connected  to  a 
supply  of  oxygen.  The  apparatus  is  supported  in  a  clamp  on  a 
retort  stand,  not  shown  in  the  figure.  A  little  plug  of  cotton  wool 
should  be  pushed  down  to  the  bottom  of  the  chimney,  so  as  to  cover 
the  open  end  of  the  narrow  tube  through  which  the  oxygen  enters  : 
this  distributes  the  oxygen  all  round  the  centre  pipe. 


1 82  Nitrogen  and  its  Compounds. 

First  regulate  the  little  gas  flame,  so  that  a  gentle  stream  of 
ammonia  escapes  up  the  centre  tube.  Then  bring  a  lighted  taper 
to  the  end  of  this  tube  and  again  notice  the  appearance  of  flame. 

Now  gently  admit  oxygen  through 
the  narrow  tube,  still  holding  the 
taper  to  the  escaping  jet  of  am- 
monia. As  soon  as  a  little  oxygen 
reaches  the  top  of  the  chimney  the 
ammonia  will  light,  and  will  con- 
tinue burning  without  the  taper, 
giving  a  curious  yellow-brown 
coloured  flame.  Now  stop  the 
oxygen,  and  notice  the  ammonia 
flame  gradually  languish  and  go 
out. 

Combination  of  Ammo- 
nia with  Acids. 

Experiment  148.— -Pour  a  little 
dilute  hydrochloric  acid,  dilute  ni- 
tric acid,  and  dilute  sulphuric  acid 
into  three  separate  little  beakers, 
and  add  two  or  three  drops  of  lit- 
mus solution  to  each.  Now  heat 

some  strong  ammonia  solution  in  a  flask,  and  pass  a  stream  of  the 
gas  into  each  of  the  acids  until  the  red  colour  of  the  litmus  changes 
to  blue,  and  then  evaporate  each  solution  to  dryness  in  separate 
dishes  heated  gently  over  small  rose  burners.  In  each  dish  a 
residue  is  left,  which  but  for  the  slight  colour  due  to  the  litmus 
would  be  white.  These  residues  are  salts  of  ammonia ;  they 
consist  of  ammonium  chloride  (sometimes  known  by  its  ancient 
name  of  sal-ammoniac}^  ammonium  nitrate,  and  ammonium  sulphate. 

Ammonium  chloride  and  ammonium  sulphate  are  com- 
mercially obtained  by  driving  off  the  ammonia  from  the 
"  ammoniacal  liquor  "  of  the  gas-works,  and  passing  the  gas 
into  either  hydrochloric  or  sulphuric  acid,  just  as  in  Exp.  148. 

The  equations  which  represent  the  combination  of  am- 
monia with  these  three  acids  are  as  follows : — 

(1)  NH3  4-  HC1     =  NH4C1. 

(2)  NH3  +  HN03  =  NH4N03. 

(3)  2NH3  +  H2S04  =  (NH4)2S04. 


Ammonia.  183 

How  to  get  Ammonia  out  of  its  Salts. 

Experiment  149. — Heat  a  small  quantity  of  the  ammonium 
chloride  obtained  in  Exp.  148,  in  a  dry  test-tube.  Note  that  the 
salt  sublimes.  Apply  a  lighted  taper  to  the  mouth  of  the  tube, 
and  see  that  it  shows  no  signs  of  a  flame  of  ammonia.  Smell  the 
tube,  and  hold  reddened  litmus  paper  to  it,  and  observe  not  even 
a  trace  of  ammonia  is  given  off.  Therefore  we  cannot  obtain 
ammonia  from  this  compound  simply  by  heating  it. 

Experiment  150. — Heat  a  little  ammonium  nitrate  in  the  same 
way.  Note  at  once  a  great  difference  in  the  behaviour  of  this  salt. 
It  melts  ;  the  chloride  did  not.  Presently  it  effervesces  ;  evidently 
gas  is  coming  off.  Is  this  gas  ammonia?  Test  with  litmus  ;  note 
no  blue  effect.  Smell  the  gas ;  there  is  no  smell  of  ammonia. 
Bring  a  lighted  taper  into  the  gas  ;  notice  that  the  gas  behaves 
like  oxygen.  Test  with  a  glowing  splint  of  wood  ;  it  relights  in  this 
gas.  Can  it  be  oxygen  ?  We  must  examine  this  point  later  ;  in  the 
mean  time,  the  experiment  shows  that  we  do  not  get  ammonia  by 
simply  heating  this  salt. 

Experiment  151. — Heat  a  small  quantity  of  ammonium  sulphate 
in  a  similar  manner.  Notice  that,  as  with  the  nitrate,  this  salt 
also  melts  and  gives  off  gas.  Test  with  litmus  and  turmeric,  and 
observe  alkalinity.  Smell  the  gas,  and  note  that  when  ammonium 
sulphate  is  heated  alone,  ammonia  is  given  off. 

Experiment  152. — Now  take  a  small  quantity  of  each  of  the 
three  ammonium  salts  in  three  test-tubes,  and  add  to  each  about 
the  same  quantity  of  powdered  lime  and  apply  heat  gently.  Notice 
that  in  each  case  ammonia  is  given  off,  as  indicated  by  litmus  or 
turmeric  paper,  as  well  as  by  the  smell. 

These  four  experiments  show  that  it  is  only  certain 
ammonium  salts  which  give  off  ammonia  when  heated  alone, 
but  that  all  of  them  yield  ammonia  when  heated  with  lime. 
We  can,  therefore,  use  this  latter  method  as  a  test  as  to 
whether  a  particular  salt  is  an  ammonium  compound  or  not. 
In  practice,  when  it  is  required  to  make  such  a  test,  we 
usually  employ  a  solution  of  sodium  hydroxide  (caustic  soda) 
instead  of  the  lime,  as  it  is  rather  more  convenient  to  use. 

Experiment  153. — Add  a  little  caustic  soda  solution  to  a  small 
quantity  of  say  ammonium  chloride  in  a  test-tube,  and  gently  warm 


184  Nitrogen  and  its  Compounds. 

the  mixture.  Smell  the  escaping  gas,  and  hold  moistened  litmus 
or  turmeric  paper  to  the  mouth  of  the  tube.  Note  abundance  of 
ammonia. 

Ammonia  Solution  is  made  by  dissolving  ammonia  gas 
in  water. 

Experiment  154. — Place  some  ammonium  chloride  in  a  flask, 
and  add  about  twice  as  much  dry  slaked  lime  and  mix  the  two 
together.  Arrange  the  flask  in  connection 
with  a  bottle  containing  water  as  shown  in 
Fig.  86,  and  gently  heat  the  flask.  Air  is 
first  expelled,  which  bubbles  through  the 
water,  but  presently,  as  only  pure  ammonia 
escapes  from  the  flask,  the  whole  of  the  gas 
is  absorbed  by  the  water,  no  bubbles  pass- 
ing through  the  water.  Notice  that,  after 
a  while,  the  liquid  in  the  bottle  begins  to 
get  perceptibly  warm  to  the  hand.  If,  there- 
fore, we  wish  to  make  a  very  strong  solu- 
tion of  the  gas,  we  must  prevent  this  by 
immersing  the  bottle  in  cold  water,  for  we 
have  learnt  by  Exp.  145,  that  by  warming  a 
solution  of  ammonia  the  gas  is  expelled 
again.  [Note.  By  arranging  the  apparatus 
so  that  the  tube  leading  into  the  water  is  a 
good  length,  there  is  no  fear  of  the  liquid 
in  the  bottle  being  sucked  back  into  the 

flask  ;  because  if  it  begins  to  ascend  far  up  the  tube  (which  might 
happen  if  the  evolution  of  gas  were  interrupted  by  a  draught 
blowing  the  flame  away  for  a  moment),  then  air  is  at  once  drawn 
into  the  flask  through  the  small  quantity  of  mercury  placed  in 
the  bend  of  the  funnel  tube.] 

The  chemical  changes  taking  place  when  ammonium 
chloride  is  heated  with  caustic  soda  and  with  slaked  lime,  are 
expressed  by  the  following  equations  : — 

NH4C1  -f  NaHO    =  NH3    +  NaCl  +  H2O 
and  2NH4C1  +  CaH2O2  -  2NH3  +  CaCl2  +  2H2O. 

The  nature  of  the  change  is  the  same  in  both  cases,  which 


The  Composition  of  Ammonia.  185 

will  be  more  easily  understood  if  the  equations  be  dissected 
in  the  following  way.     (i)  With  the  sodium  hydroxide  — 

NH3 

H    +    JHO    =    H20 
Cl         (Na  NaCl 


(NH 

\ 

V 


NH4C1+  NaHO  -  NaCl  +  H2O  +  NH8 

and  (2)  with  slaked  lime  (or  calcium  hydroxide) — 
NH3  NH3  NH3 

H  +  (HO  HO  =  H20  H2O 
Cl      1      Ca          CaCl. 


(NH, 

H 

I    Cl 


2NH4C1        -f    CaH2O2  =  CaCl2 

The  Composition  of  Ammonia. — We  have  learnt  that 
hydrogen  and  oxygen  unite  together  very  readily,  forming 
water;  also  that  hydrogen  and  chlorine  combine  with  great 
ease,  giving  hydrochloric  acid,  so  that  we  were  able  to  gain 
knowledge  of  the  composition  of  both  water  and  hydrochloric 
acid  by  synthesis.  But  hydrogen  only  combines  directly  with 
nitrogen  with  great  difficulty,  and  even  then  under  exceptional 
conditions,  therefore  we  cannot  find  the  composition  of 
ammonia  by  the  direct  union  of  its  elements.  We  can,  how- 
ever, decompose  ammonia,  and  find  the  proportional  volumes 
of  its  constituents.  We  do  this  by  means  of  chlorine,  which 
has  been  shown  (p.  176)  to  be  able  to  decompose  this  com- 
pound, combining  with  the  hydrogen  and  letting  the  nitrogen 
go  free. 

Experiment  155. — Take  a  long  glass  tube,  closed  up  at  one  end, 
and  divide  it  into  three  equal  divisions  with  indiarubber  bands. 
Fit  to  the  tube  a  cork  carrying  a  small  stoppered  dropping  funnel, 
as  shown  in  Fig.  87.  Collect  this  long  tube  full  of  chlorine  (using 
strong  brine  in  the  pneumatic  trough)  and  insert  the  cork.  Pour 
a  little  strong  ammonia  solution  into  the  funnel,  and  allow  it  slowly 
to  enter  the  tube,  one  drop  at  a  time.  The  first  few  drops  take  fire 
as  they  go  in  (as  in  Exp.  143). 

When  about  a  dozen  drops  have  been  let  in,  the  action  is  finished. 
Fill  up  the  funnel  with  dilute  sulphuric  acid,  so  as  to  neutralize  the 


1 86 


Nitrogen  and  its  Compounds. 


excess  of  ammonia,  and  fit  a  bent  tube  into  the  mouth  of  it,  and  let 
the  long  end  dip  into  a  beaker  of  water  as  arranged  in  the  figure. 
Then  open  the  tap.  Water  at  once  runs 
in,  showing  that  some  gas  has  disappeared, 
and  it  will  fill  up  exactly  two  of  the  measures, 
leaving  one  measure  of  gas. 

Test  this  gas  by  removing  the  cork  and 
dipping  a  lighted  taper  into  the  tube.  The 
gas  is  nitrogen.  We  therefore  have  got 
one  measure  of  nitrogen,  which  has  been 
expelled  from  combination  with  just  so 
much  hydrogen  as  there  was  chlorine  ori- 
ginally present  in  the  tube,  namely,  three 
measures  (because  chlorine  and  hydrogen 
combine  in  equal  volumes.  See  Exp.  103), 
Therefore,  in  ammonia,  the  nitrogen  and 
hydrogen  are  combined  in  the  proportion 
one  volume  of  nitrogen  to  three  volumes  of 
hydrogen. 

In  order  to  prove  that  the  formula 
is  NH3  and  not  N2H6  (both  of  which 
contain  the  two  elements   in  the  same 
FIG.  87.  relative  proportion),   it  is  necessary  to 

find  out  the  density  of  the  gas  by  weighing  a  known  volume 
of  it.  When  this  is  done,  it  is  found  that  ammonia  gas  is 
8*5  times  as  heavy  as  hydrogen.  Its  molecular  weight,  there- 
fore, is  twice  this,  namely,  17.  This  proves  that  the  composi- 
tion is  expressed  by  the  formula  NH3. 

N  =  14.     3H  =  3.     .'.  14  +  3  =  17. 

EPITOME. 

Ammonia  is  formed  by  the  putrefaction  of  organic  matter  con- 
taining nitrogen,  such  as  manure.  Minute  quantities  of  it  are 
found  in  the  air. 

Ammonia  is  produced  during  the  destructive  distillation  of  coal 
(as  in  the  manufacture  of  ordinary  coal  gas),  when  it  collects  in  the 
watery  liquid  known  as  the  "  ammoniac al  liquor." 

Ammonia  is  obtained  from  its  salts  by  heating  them  with  slaked 
lime  or  with  sodium  hydroxide  (caustic  soda) 

Ammonia  is  a  colourless  gas,  with  a  powerful  pungent  smell, 


Ammonia.  187 

and  a  strong  alkaline  reaction.  It  is  very  soluble  in  water,  and 
therefore  cannot  be  collected  in  the  ordinary  way.  One  litre  of 
water  at  the  common  temperature  dissolves  about  800  litres  of 
ammonia  gas,  while  at  o°  it  will  absorb  as  much  as  1148  litres.  All 
the  gas  is  driven  out  of  solution  when  the  liquid  is  heated. 

Ammonia  will  not  burn  in  air,  but  it  will  burn  in  oxygen.  It  will 
not  support  the  combustion  of  a  taper.  Ammonia  is  a  light  gas  ; 
it  is  just  over  one  half  as  heavy  as  air,  and  is  8*5  times  as  heavy  as 
hydrogen. 

Ammonia  is  easily  condensed  to  the  liquid  state.  At  the  ordinary 
temperatures,  about  7  atmospheres  pressure  will  squeeze  this  gas 
into  the  liquid  form.  And  again,  if  the  gas,  without  being 
squeezed,  is  simply  cooled  below  -  34°  C.  then  it  also  turns  into 
liquid  ammonia. 

[Note. — Liquid  ammonia  is  not  the  same  as  a  solution  of 
ammonia  in  water.] 

Liquid  ammonia  boils  at  —  337°,  and  has  been  largely  used  for 
the  artificial  production  of  ice. 

When  passed  through  a  red-hot  tube,  or  submitted  to  electric 
sparks,  ammonia  gas  is  decomposed  into  nitrogen  and  hydrogen. 
When  a  measured  volume  of  the  gas  is  thus  decomposed,  the  volume 
is  doubled;  that  is  to  say,  two  volumes  of  ammonia  give  one  volume 
of  nitrogen  and  three  volumes  of  hydrogen. 

The  volume  composition  of  ammonia  is  also  established  by 
decomposing  it  with  chlorine,  when  it  is  found  that  three  volumes 
of  chlorine  take  three  volumes  of  hydrogen  to  form  hydrochloric 
acid,  and  leave  one  volume  of  nitrogen. 


CHAPTER   XX. 

NITROGEN   AND    ITS   COMPOUNDS    (continued). 

Nitric  Acid,  HNO3. — This  substance  is  one  of  the  most 
important  of  all  the  nitrogen  compounds,  and  from  it,  either 
directly  or  indirectly,  we  obtain  all  the  other  compounds  of 
nitrogen  which  it  will  be  necessary  for  us  to  study.  Nitric 
acid  is  composed  of  nitrogen,  hydrogen,  and  oxygen,  but  we 
do  not  make  it  by  causing  these  three  elements  to  unite  directly. 
It  can  be  produced  synthetically,  however,  for  if  we  add  a  little 
nitrogen  to  a  mixture  of  oxygen  and  hydrogen,  and  ignite  the 
mixture,  the  water  which  results  from  the  union  of  the  hydrogen 
and  oxygen  will  be  found  to  be  acid,  from  the  presence  of  a 
little  nitric  acid.  Cavendish  noticed  this  when  making  his 
experiments  on  the  composition  of  water  (p.  73),  and  was  for 
a  long  time  puzzled  to  account  for  the  acidity  of  the  water  he 
got,  but  discovered  that  nitrogen  from  the  air  had  found  its 
way  into  the  apparatus. 

If  a  rapid  stream  of  electric  sparks  is  passed  between 
platinum  wires  in  a  confined  space  of  air,  the  oxygen  and 
nitrogen  begin  to  combine,  forming  an  oxide  of  nitrogen.  This 
gas  has  a  brownish  colour,  and  if  a  little  water  be  then  added 
the  brown  gas  disappears ;  it  dissolves  in  the  water  and  yields 
nitric  acid.  In  this  manner  nitric  acid  is  produced  during 
thunderstorms.  The  lightning  flashes  (which  are  simply 
enormous  electric  sparks)  passing  through  the  air  cause  the 
combination  of  some  of  the  nitrogen  and  oxygen,  and  the  oxide 
so  formed  is  washed  out  by  the  rain. 

Preparation. — Nitric  acid  is  made  from  either  potassium 


Nitric  Acid. 


1 89 


FIG.  88. 


nitrate  (nitre  or  saltpetre}  or  from  sodium  nitrate  (Chili 
saltpetre). 

Experiment  156. — Place  50  grams  of  nitre  in  a  glass  retort 
with  a  stopper,  and  pour 
upon  it  the  same  weight 
of  strong  sulphuric  acid. 
Gently  heat  the  mixture 
and  collect  the  distillate  in 
a  clean  flask,  which  is  kept 
cool  by  being  placed  in 
a  dish  of  cold  water,  as 
shown  in  Fig.  88.  Now  and 
then  turn  the  flask  round, 
so  as  to  cool  it  all  over. 
Notice  that  reddish  fumes 
appear  in  the  retort,  and 
that  a  liquid  collects  in  the  flask  which  has  a  pale  yellowish  colour. 

The  equation  which  takes  place  in  this  experiment  is  the 
following — 

KNO3  +  H2SO4  -  HKS04  +  HNO3. 

The  potassium  in  the  nitre  changes  places  with  one-half  of 
the  hydrogen  in  the  sulphuric  acid,  giving  a  salt  called  hydrogen 
potassium  sulphate,  and  nitric  acid.  The  salt  remains  behind 
in  the  retort. 

If  sodium  nitrate  had  been  used  the  reaction  would  have 
been  quite  similar. 

The  manufacturer  of  nitric  acid  always  employs  the  Chili 
saltpetre,  because  it  is  a  cheaper  article  than  the  potassium 
salt.  He  also  carries  on  the  operation  at  a  higher  temperature, 
using  large  cast-iron  vessels  in  the  place  of  a  glass  retort,  which 
enables  him  to  get  all  the  hydrogen  in  his  sulphuric  acid  ex- 
changed for  sodium,  according  to  the  equation — 

2NaNO3  +  H2SO4  =  Na2SO4  +  2HNO3. 

The  same  quantity  of  sulphuric  acid,  therefore,  is  made  to 
decompose  twice  as  much  of  the  nitrate  as  in  the  laboratory 
experiment.  At  the  high  temperature,  however,  required  to 


Nitrogen  and  its  Compounds. 

completely  bring  about  this  second  reaction,  a  little  of  the  nitric 
acid  itself  is  decomposed,  and  therefore  wasted. 

Properties  of  Nitric  Acid.— When  quite  pure,  the  acid 
is  colourless.  The  sample  obtained  in  Exp.  156  is  coloured 
slightly  yellow,  because  it  has  dissolved  some  of  the  coloured 
fumes  which  were  produced, 

Nitric  acid  fumes  strongly  when  exposed  to  moist  air.  It 
can  be  mixed  with  water  in  any  proportions.  The  strong  acid 
is  highly  corrosive,  and  must  be  handled  with  great  care  ;  a  few 
drops  spilt  upon  the  skin  will  cause  bad  wounds.  Even  when 
moderately  diluted  with  water  it  will  burn  the  clothes,  and  stain 
the  skin  yellow.  If  strong  nitric  acid  is  boiled,  it  begins  to 


FIG.  89. 

decompose,  giving  oxygen  and  nitrogen  peroxide  (a  reddish- 
brown  gas).  When  heated  strongly  this  decomposition  is  very 
rapid. 

Experiment  157. — Arrange  a  tobacco-pipe  as  shown  in  the 
figure,  the  mouthpiece  just  dipping  beneath  the  water  in  the  trough. 
When  the  stem  is  red-hot,  pour  a  few  drops  of  strong  nitric  acid 
into  the  bowl.  As  the  acid  passes  the  heated  place  it  is  decomposed 
into  the  two  gases,  oxygen  and  nitrogen  peroxide.  The  latter  gas, 
however,  is  very  soluble  in  water,  and  therefore  dissolves  in  the 
trough,  while  oxygen  alone  is  collected. 

Notice  that,  as  the  bubbles  first  appear  in  the  water,  they  have 
a  dark  reddish  colour,  while  the  gas  which  actually  collects  is 
colourless.  Test  the  gas  for  oxygen.  The  equation  is  the 
following — 

2HNO3  =  H2O  +  2NO2  +  O. 

Nitric  acid  is  therefore  a  powerful  oxidizing  material,  and 
it  acts  on  many  substances  with  great  energy. 


Nitric  Acid.  191 

Experiment  158. — Gently  heat  a  small  quantity  of  sawdust  in  a 
small  porcelain  dish  or  crucible,  until  the  wood  has  just  become 
charred.  Then  cautiously  let  two  or  three  drops  of  strong  nitric 
acid  fall  upon  the  charred  mass,  and  notice  that  it  instantly  takes 
fire.  The  charcoal  burns  at  the  expense  of  the  oxygen  supplied  by 
the  nitric  acid,  while  fumes  of  the  brown  gas  are  produced  at  the 
same  time. 

Experiment  159. — Add  a  little  powdered  sulphur  to  a  small 
quantity  of  strong  nitric  acid  in  a  test-tube,  and  gently  heat  the 
mixture.  Notice  that  the  brown-coloured  oxide  of  nitrogen  is  again 
given  off  in  quantity.  The  acid  is  therefore  being  reduced;  in  other 
words,  it  is  giving  up  some  of  its  oxygen  to  the  sulphur,  which  is 
gradually  being  converted  into  sulphuric  acid.  Let  the  action  con- 
tinue for  a  few  minutes,  and  then  test  the  liquid  for  sulphuric  acid 
in  the  following  way — 

Add  one  or  two  drops  of  sulphuric  acid  to  some  water  in  a  test- 
tube,  and  then  add  some  solution  of  barium  chloride.  Notice  a 
white  precipitate.  The  reaction  here  is 

H2SO4  +  BaCl2  =  BaSO4  +  2HC1. 

The  barium  and  hydrogen  change  places,  and  barium  sulphate  is 
produced.  This  Substance,  being  insoluble  in  water,  separates  out 
as  a  white  solid.  There  are,  however,  other  things  which  would  give 
a  white  precipitate  with  barium  chloride,  but  barium  sulphate  may 
be  distinguished  from  the  other  precipitates  by  the  fact  that  it 
cannot  be  dissolved  by  acids.  Therefore,  as  a  confirmation  of  the 
test,  pour  half  the  white  precipitate  into  a  second  test-tube,  and  to 
one  portion  add  some  strong  hydrochloric  acid,  and  to  the  other 
some  nitric  acid,  and  note  that  the  precipitate  is  riot  dissolved  in 
either  case. 

Now  apply  this  test  in  the  case  of  the  nitric  acid  in  which 
sulphur  has  been  heated.  Dilute  the  acid  with  water,  and  add  a 
few  drops  of  barium  chloride.  A  white  precipitate  proves  the 
presence  of  sulphuric  acid,  which  requires  no  confirmatory  test,  as 
the  liquid  already  contains  nitric  acid. 

Besides  oxidizing  such  elements  as  carbon,  sulphur,  phos- 
phorus, iodine,  etc.,  it  attacks  a  great  many  metals,  converting 
them  into  compounds  which  contain  oxygen,  and  are,  therefore, 
oxidation  products. 

Experiment  1 60.— Drop  a  few  fragments  of  copper  wire  or 
foil  into  a  little  nitric  acid  in  a  test-tube.  Notice  that  violent 


192  Nitrogen  and  its  Compounds. 

action  at  once  sets  in.  Torrents  of  the  red  gas  are  given  off,  and 
the  copper  quickly  disappears,  while  the  liquid  becomes  blue.  The 
copper  is  converted  into  copper  nitrate  (a  blue  salt),  which  remains 
dissolved  in  the  liquid.  A  number  of  other-  metals  behave  in  a 
similar  way,  such  as  silver,  mercury,  iron,  lead,  zinc.  They  are 
converted  into  nitrates,  while  the  nitric  acid  is  reduced  to  one  of 
the  oxides  of  nitrogen. 

We  may  suppose  that  the  first  action  of  the  acid  on  such  metals 
is  that  the  metal  changes  place  with  the  hydrogen,  thus — 

Cu+  2HN03  =  H2  +  Cu(N03)2, 

and  that  immediately  this  nascent 1  hydrogen  attacks  another 
portion  of  nitric  acid,  taking  away  some  of  its  oxygen,  whereby 
water  is  formed  and  an  oxide  of  nitrogen  left.  Which  oxide  of 
nitrogen  is  left  depends  on  a  number  of  circumstances,  for  hydrogen 
in  this  state  can  gradually  take  all  the  oxygen  from  nitric  acid, 
leaving  at  last  only  nitrogen.  In  no  case  by  the  action  of  nitric 
acid  on  metals  is  any  hydrogen  given  off  so  as  to  be  collected.  On 
account  of  its  powerful  action  in  dissolving  metals,  nitric  acid  used 
to  be  called  aqua-fortis  (the  strong  water)  ;  but  even  this  acid 
cannot  dissolve  the  "  noble  "  metals,  gold  and  platinum. 

Aqua  Regia. — Both  nitric  acid  and  hydrochloric  acid 
are  without  any  action  on  gold  or  platinum,  but  if  the  two 
acids  are  mixed  together,  then  the  mixture  will  easily  dissolve 
either  of  these  metals. 

Experiment  161. — Place  a  gold  leaf  in  a  wide  test-tube  and  pour 
strong  nitric  acid  upon  it.  The  leaf  breaks  up  into  small  particles, 
but  notice  that  it  is  not  dissolved,  and  may  be  left  a  long  time  in 
the  acid.  Do  the  same  with  hydrochloric  acid  in  another  test- 
tube,  and  notice  that,  in  like  manner,  the  gold  leaf  is  not  acted 
upon.  Now  mix  the  contents  of  the  two  tubes,  and  in  a  few 
moments  the  gold  will  be  entirely  dissolved. 

The  mixture  of  these  two  acids  is  known  as  aqua  regia  (the 
royal  water),  just  because  it  is  able  to  dissolve  the  noble 
metals.  When  metals  dissolve  in  aqua  regia,  it  is  always  the 
chloride,  and  not  the  nitrate  of  the  metal,  that  is  formed. 

1  Nascent  means  just  born ;  and  elements  at  the  moment  of  their 
liberation  from  combination,  are  said  to  be  in  the  nascent  state  ;  at  this 
particular  moment  the  element  is  much  more  chemically  active. 


Nitrates.  193 

Impurities  in  Nitric  Acid. — Common  nitric  acid 
generally  contains  a  number  of  impurities,  the  most  usual  of 
which  are  sulphuric  acid  (derived  from  the  sulphuric  acid 
used  in  its  manufacture)  and  iron  (from  the  vessels  in  which 
the  acid  is  made).  The  sulphuric  acid  may  be  tested  for  by 
the  method  described  in  Exp.  159.  The  presence  of  iron  in 
the  acid  may  be  found  out  by  the  following  test. 

Experiment  162. — Dilute  a  little  of  the  suspected  acid,  and  add 
to  it  a  few  drops  of  a  solution  of  potassium  ferrocyanide  (yellow 
prussiate  of  potash}.  If  much  iron  is  present,  a  deep  blue  precipi- 
tate will  be  seen  ("  Prussian  blue  ")  ;  but  if  there  is  only  a  little  of 
this  impurity  in  the  acid,  the  ferrocyanide  will  only  produce  a 
bluish  or  greenish  colour. 

Nitrates. — The  salts  which  nitric  acid  forms,  when  its 
hydrogen  is  replaced  by  metals,  are  called  nitrates ;  such  as 
potassium  nitrate,  copper  nitrate,  etc. 

Like  hydrochloric  acid,  nitric  acid  has  only  one  atom  of 
hydrogen  in  it ;  it  is  on  this  account  called  a  mono-basic  acid. 
Nitrates  are  not  only  produced  by  dissolving  metals  in  the 
acid,  but  also  by  acting  on  metallic  oxides,  hydroxides,  or 
carbonates,  with  nitric  acid. 

Experiment  163. — Place  in  three  separate  dishes  a  little  copper 
oxide,  potassium  hydroxide,  and  sodium  carbonate.  Add  a  little 
strong  nitric  acid  to  the  copper  oxide  and  gently  warm  it.  Notice 
that  the  oxide  dissolves,  giving  the  same  blue  solution  of  copper 
nitrate  as  in  Exp.  160,  but  that  no  brown  gas  is  given  off. 

Dilute  some  nitric  acid,  and  add  it  gradually  to  the  potassium 
hydroxide  and  the  sodium  carbonate  until  each  is  dissolved.  Now 
slowly  evaporate  all  three  solutions  down  to  dryness.  Blue  crystals 
of  copper  nitrate,  and  white  crystals  of  potassium  and  sodium 
nitrates  will  be  obtained.  The  following  are  the  three  equations 
for  the  formation  of  these  nitrates — 

(1)  CuO  +  2HNO3  =  Cu(NO3)2  +  H2O 

(2)  KHO  +  HNO3  =  KNO3  +  H2O 

(3)  Na2CO3  +  2HNO3  =  2NaNO3  +  CO2  +  H2O. 

Experiment  164. — Spread  a  drop  or  two  of  the  solutions  of 
sodium  nitrate  and  potassium  nitrate  obtained  in  the  last  experi- 
ment, upon  separate  pieces  of  clean,  flat  glass,  and  allow  the 

O 


194  Nitrogen  and  its  Compounds 

solution  to  evaporate  by  itself.  Carefully  compare  the  crystals  of 
the  one  salt  with  those  of  the  other,  using,  if  necessary,  a  pocket 
lens.  The  shapes  of  the  crystals  are  quite  different.  Those  of 
potassium  nitrate  are  long  thin  prisms,  while  the  sodium  nitrate 
crystals  are  more  like  little  cubes  ;  although  not  exactly  cubes. 
[Sodium  nitrate  is  sometimes  called  cubical  nitre  for  this  reason.] 

All  nitrates,  when  strongly  heated,  are  decomposed,  and, 
in  most  cases,  they  give  off  oxygen.  They  are,  therefore,  like 
nitric  acid  itself,  powerful  oxidizing  agents,  and  will  readily 
give  up  oxygen  to  substances  capable  of  taking  it. 

Experiment  165. — Heat  a  few  crystals  of  potassium  nitrate  in  a 
test-tube.  They  first  melt,  and  presently  give  off  gas.  Now  drop 
into  the  test-tube  a  fragment  of  charcoal,  about  the  size  of  a  grain 
of  corn.  The  charcoal  takes  fire  when  it  touches  the  hot  nitre,  and 
the  little  fragment  dances  about  on  the  molten  salt.  Next  drop  in 
a  piece  of  sulphur  about  the  same  size,  and  note  how  readily  it 
burns  in  contact  with  the  melted  nitre. 

These  three  substances,  nitre,  charcoal,  and  sulphur,  are 
what  gunpowder  is  composed  of.  The  nitre  affords  a  supply 
of  oxygen  sufficient  to  burn  up  the  carbon  and  sulphur,  so  that 
these  materials  can  burn  in  places  where  they  cannot  get 
oxygen  from  the  air,  such  as  in  the  breech  of  a  gun,  or  even 
underneath  water. 

Experiment  166.- — Weigh  out  15  grams  of  finely  powdered  nitre, 
3  grams  of  charcoal  (also  finely  powdered),  and  2  grams  of  flowers 
of  sulphur.  Mix  these  together  most  thoroughly,  not  with  a  pestle 
and  mortar,  but  by  putting  them  all  into  a  sieve,  and  shaking  them 
through  together  (a  sieve  may  easily  be  made  by  tying  a  piece  of 
muslin  over  the  mouth  of  a  beaker,  the  bottom  of  which,  has  been 
broken). 

Make  a  little  heap  with  a  portion  of  this  mixture,  upon  a  piece 
of  wood,  and  set  fire  to  it  with  a  match  or  taper.  Notice  how  it 
burns— not  so  suddenly  as  proper  gunpowder  burns,  because  the 
manufacturer  is  able  to  get  the  ingredients  much  more  intimately 
mixed  than  in  this  case. 

Experiment  167.— Make  a  little  paper  tube  about  7  or  8  centi- 
metres long  (3  inches)  by  rolling  a  piece  of  writing  paper  closely 
round  a  lead-pencil  about  half  a  dozen  turns.  Fasten  the  edge  down 
with  a  little  strong  gum.  Slip  the  pencil  out,  and  stop  up  one  end 


Tests  for  Nitrates.  195 

of  the  paper  tube  with  a  tiny  cork  and  a  little  sealing-wax.  Now 
proceed  to 'fill  the  tube  with  the  mixture  prepared  in  the  last 
experiment,  putting  in  a  little  at  a  time,  and  packing  it  tightly 
down,  using  the  lead-pencil  as  a  ram-rod.  It  is  now  much  the 
same  thing  as  an  ordinary  "  squib,"  but  without  the  "  bang." 

Now  light  the  open  end,  and  while  it  is  burning  plunge  it  under 
water,  notice  that  it  continues  burning,  and  gives  off  a  quantity  of 
gas  which  bubbles  up  through  the  water. 

Tests  for  Nitrates. — All  nitrates  are  soluble  in  water, 
and  the  presence  of  such  a  salt  in  a  solution  can  easily  be 
detected  by  either  of  the  two  following  tests. 

Experiment  168. — Dissolve  a  crystal  of  potassium  nitrate  in  a 
little  water,  and  to  a  part  of  the  solution  so  obtained  add  some  strong 
sulphuric  acid.  The  nitrate  is  converted  into  free  nitric  acid  (just 
as  in  the  preparation  of  nitric  acid.  Exp.  1 56).  Then  drop  into  the 
mixture  a  fragment  or  two  of  metallic  copper.  As  we  have  learnt, 
nitric  acid  acts  on  copper,  and  red  fumes  are  evolved  ;  so  that  if 
the  liquid  containing  the  copper  be  now  gently  warmed,  red  fumes 
will  be  given  off,  which  proves  that  a  nitrate  was  originally  present. 

The  second  test  is  more  delicate,  and  will  detect  much  smaller 
quantities  of  a  nitrate.  Take  another  portion  of  the  solution  con- 
taining the  nitrate,  and,  as  before,  add  sulphuric  acid.  Now  pour 
gently  down  into  the  test-tube  a  solution  of  ferrous  sulphate,  so  that 
this  solution  floats  on  the  top  of  the  other.  Notice  that  where  the 
two  liquids  meet  in  the  tube,  there  is  a  dark  brown  layer  formed  ; 
which,  if  the  ferrous  sulphate  has  been  added  quite  carefully  so  as 
not  to  get  mixed  with  the  other  solution,  will  appear  as  a  well 
defined  ring. 

[The  nitric  acid  produced  by  the  addition  of  sulphuric  acid  to 
the  nitrate,  gives  up  some  of  its  oxygen  to  the  ferrous  sulphate,  and 
a  small  quantity  of  one  of  the  oxides  of  nitrogen  is  formed.  This 
dissolves  in  a  further  portion  of  ferrous  sulphate,  forming  a  dark 
brown  compound.] 

EPITOME. 

Nitric  acid  (aqua  fortis)  is  made  by  the  action  of  sulphuric 
acid  on  potassium  or  sodium  nitrate.  On  the  manufacturing  scale 
sodium  nitrate  is  used  (Chili  saltpetre),  as  this  salt  is  cheaper. 
Nitric  acid  is  produced  when  electric  sparks  pass  through  air  in 
the  presence  of  water,  hence  it  is  formed  during  thunderstorms,  and 
gets  washed  into  the  ground  with  the  rain. 


195  Nitrogen  and  its  Compounds, 

Pure  nitric  acid  is  a  colourless  fuming  corrosive  liquid.  It  begins 
to  decompose  when  boiled,  giving  red  fumes  of  nitrogen  peroxide, 
and  oxygen. 

Nitric  acid  attacks  most  metals,  but  has  no  action  on  gold  or 
platinum.  The  result  of  the  action  of  nitric  acid  on  metals  is 
either  a  nitrate  or  an  oxide  of  the  metal,  and  one  or  more  of  the 
oxides  of  nitrogen  :  but  hydrogen  is  never  evolved. 

We  can  at  once  distinguish  gold  from  any  imitations  of  this 
metal,  by  touching  the  surface  with  a  drop  of  nitric  acid.  If  there 
is  no  action,  the  material  is  gold. 

Nitric  acid  is  a  powerful  oxidizing  substance,  and  converts  the 
elements  sulphur,  phosphorus,  and  iodine  into  sulphuric,  phosphoric, 
and  iodic  acids  respectively,  with  evolution  of  oxides  of  nitrogen. 

The  salts  of  nitric  acid  are  nitrates.  Potassium  nitrate  (nitre, 
saltpetre]  is  one  of  the  most  important.  It  is  a  constituent  of  gun- 
powder. When  heated,  nitrates  decompose  and  give  up  oxygen  ; 
sometimes  forming  first  of  all  a  nitrite,  thus — 

KNO3  =  O  +  KNO2 

Nitrates  of  certain  metals,  when  heated,  leave  an  oxide  of  the 
metal,  and  give  off  nitrogen  peroxide  and  oxygen,  thus — 

Pb(NO3)2  =  PbO  +  2NO2  -f  O. 


CHAPTER    XXI. 

OXIDES    OF    NITROGEN. 

THERE  are  five  oxides  of  nitrogen  ;  their  names  and  formulae 
are  as  follows  : — 

(r)  Nitrous  oxide  (laughing gas)         ...          ...  N2O 

(2)  Nitric  oxide  NO 

(3)  Nitrogen  trioxide N2O3 

(4)  Nitrogen  peroxide  ...          ...          ...         ...  N(X 

(5)  Nitrogen  pentoxide  N2O5 

Numbers  i,  2,  and  5  are  acidic  oxides.  The  acid  derived 
from  nitrous  oxide  (called  hyponitrous  acid)  and  that  from 
nitrogen  trioxide  (nitrous  acid)  are  both  very  unstable  com- 
pounds, and  have  never  been  obtained  in  the  free  state.  The 
acid  derived  from  nitrogen  pentoxide  is  nitric  add. 

Nitric  Oxide,  NO.— It  will  be  most  convenient  to 
study  this  oxide  first. 

When  nitric  acid  acts  on  metals,  as  explained  on  page  192, 
the  hydrogen  which  is  first  displaced  from  the  acid  by  ihe 
metal,  at  once  attacks  a  further  portion  of  nitric  acid,  depriving 
it  of  more  or  less  of  its  oxygen,  and  thereby  reducing  it  to 
one  or  other  of  the  oxides  of  nitrogen.  The  gradual  reduction 
of  nitric  acid  by  hydrogen  will  be  seen  by  the  following 
equations  :— 

(1)  2HNO3  +    2H  =  2H2O  +  2NO2 

(2)  2HN03  +    4H  =  3H20  +  N2O3 

(3)  2HN03-f    6H  =  4H,0  +  2NO 

(4)  2HNO3  -f    8H  =  sH2O  -f  N,O 

(5)  2HNO3  +  loH  =  6H20  -f  Na. 


198  Oxides  of  Nitrogen. 

The  reducing  action  may  even  go  a  step  further,  and  re- 
sult in  the  formation  of  ammonia,  thus — 

HN03  +  8H  =  3H20  +  NH3. 

Now  it  depends  partly  on  what  metal  is  being  acted  on  by 
the  nitric  acid,  partly  on  the  strength  of  the  acid  used,  partly 
on  the  temperature,  and  partly  on  the  amount  of  the  nitrate 
of  the  metal  which  is  produced  during  the  action,  which 
particular  oxide  will  be  produced. 

For  instance,  when  nitric  acid  is  poured  upon  copper,  at 
first  some  nitrogen  peroxide  is  formed,  then  nitric  oxide  is 
evolved.  After  a  time  nitrous  oxide  begins  to  come  off.  and 
later  on  nitrogen  is  formed.  Probably  at  no  single  moment 
is  any  one  oxide  alone  produced,  so  that  we  do  not  get  any 
one  quite  free  from  the  others  by  such  an  experiment.  In 
many  cases,  however,  we  know  the  particular  conditions  which 
will  give  us  the  oxide  we  want  in  a  state  which  is  sufficiently 
free  from  the  others  for  ordinary  experiment. 

Thus,  in  order  to  get  nitric  oxide,  we  use  copper  and 
nitric  acid,  doing  the  experiment  in  the  following  way. 

Experiment  169. — Place  a  quantity  of  copper  clippings  in  a  two- 
necked  bottle,  arranged  as  in  Fig.  39,  p.  45.  Pour  upon  the  copper 
a  considerable  quantity  of  a  mixture  consisting  of  one  part  of  strong 
nitric  acid  and  two  parts  of  water  (by  measure).  In  a  few  minutes 
a  brisk  action  sets  in,  and  for  some  little  time  the  gas  which  is 
evolved  consists  cMefly  of  nitric  oxide.  Collect  three  or  four  jars 
of  the  gas  as  quickly  as  possible,  and  then  empty  out  the  apparatus. 
Avoid  breathing  the  gas,  as  it  is  injurious. 

The  reaction  which  expresses  the  production  of  nitric 
oxide,  is  that  given  in  equation  No.  3  above,  where  six  atoms 
of  hydrogen  act  on  two  molecules  of  nitric  acid,  the  six 
atoms  of  hydrogen  being  displaced  from  six  molecules  of 
nitric  acid.  The  two  steps  or  stages  in  the  reaction  may  be 
expressed  by  the  two  equations — 

3Cu  +  6HN03  =  6H  +  3Cu(N03) 
2HNO3  +  6H  =  4H2O  +  2NO; 
Or  to  express  the  final  results  by  a  single  equation — 
3Cu  4-  8HNO,  =  3Cu(N03)2  +  4H2O  +  2NO. 


Nitric  Oxide.  199 

As  already  stated,  other  reactions  resulting  in  the  formation 
of  other  oxides  of  nitrogen  are  liable  to  go  on  at  the  same 
time  as  this  one  ;  hence  when  chemists  want  perfectly  pure 
nitric  oxide,  they  ma^e  it  by  other  methods. 

Properties  of  Nitric  Oxide. — The  samples  collected 
show  that  it  is  a  colourless  gas,  which  is  not  appreciably 
soluble  in  water.  Its  most  characteristic  property  is  its  be- 
haviour when  it  comes  in  contact  with  the  air. 

Experiment  170. — Take  one  of  the  jars  of  gas  out  of  the  trough, 
and  for  a  moment  uncover  its  mouth.  Notice  that  where  the  gas 
and  air  meet,  dark  red  fumes  are  formed.  [This  jar  may  be  re- 
turned to  the  trough,  as  it  will  do  for  another  experiment.] 

Which  of  the  constituents  of  the  air  is  the  cause  of  these 
red  fumes  ? 

Experiment  171. — Decant  some  of  the  gas  from  one  of  the  jars 
into  a  narrow  cylinder  filled  with  water  and  standing  in  the  trough, 
until  the  cylinder  is  about  three  parts  filled  with  gas.  Now  bubble 
up  into  this  a  little  oxygen  [either  from  another  cylinder,  or  better 
from  a  bottle  of  compressed  gas].  Notice  that  as  each  bubble  of 
oxygen  comes  into  the  nitric  oxide,  it  forms  the  same  red  gas  as 
before.  Therefore,  it  was  the  oxygen  in  the  air  which  caused  it 
in  the  former  experiment. 

Notice  that,  as  each  addition  of  oxygen  is  made,  there  is  a 
momentary  expansion,  due  to  the  heat  produced  by  the  combination 
of  the  two  gases ;  and  also  that,  after  a  moment,  the  volume 
diminishes  and  the  red  gas  disappears. 

The  compound  that  is  formed  is  nitrogen  peroxide,  NO2. 
This  is  the  red  gas,  and  it  is  a  gas  which  is  quickly  dissolved 
by  water.  If,  therefore,  the  sample  of  nitric  oxide  is  quite 
pure,  and  oxygen  is  added  cautiously  with  frequent  shaking 
up  with  the  water,  the  entire  quantity  of  gas  ought  to  dis- 
appear altogether,  being  wholly  converted  into  nitrogen  per- 
oxide, which  will  dissolve  in  the  water. 

Since  nitric  oxide  combines  with  atmospheric  oxygen  the 
moment  it  mixes  with  the  air,  of  course  we  cannot  tell  whether 
the  gas  has  any  smell,  as  before  we  could  smell  it,  it  is  no 
longer  nitric  oxide,  but  nitrogen  peroxide;  and  this  gas  has 
a  most  unpleasant  smell  and  is  poisonous. 


2OO  Oxides  of  Nitrogen. 

Next  try  the  action  of  combustibles  on  nitric  oxide. 

Experiment  172. — Introduce  a  lighted  taper  ;  note  that  the  gas 
does  not  burn,  and  that  it  at  once  puts  out  the  taper. 

Place  a  small  piece  of  phosphorus  in  a  deflagrating  spoon  ;  set 
fire  to  it,  and  before  it  has  time  to  burn  up,  plunge  it  into  a  jar  of 
the  gas.  The  phosphorus  will  be  extinguished.  Now  withdraw 
it,  and  allow  it  to  burn  up  brightly  before  putting  it  into  the  gas  ; 
note  that  now  it  burns  in  the  gas  with  a  bright  flame,  and  forms 
white  fumes,  like  those  produced  when  it  burns  in  air  or  oxygen. 

This  shows  that  nitric  oxide  itself  does  not  support  the 
combustion  of  common  combustibles,  but  that  if  the  burning 
substance  is  hot  enough  to  decompose  the  gas  into  nitrogen 
and  oxygen,  then  combustion  goes  on  in  the  oxygen  so 
liberated.  Therefore,  the  product  of  burning  is  the  same  as 
in  the  air,  but  as  there  is  a  larger  proportion  of  oxygen  in 
nitric  oxide  than  in  air,  the  combustion  is  more  rapid  and 
brilliant. 

The  composition  of  nitric  oxide  is  found  by  taking  a 
measured  volume  of  the  gas,  and  strongly  heating  in  it  some 
metal,  such  as  iron,  which  can  combine  with  the  oxygen,  and 
leave  the  nitrogen.  When  this  is  done,  it  is  found  that  the 
volume  of  the  nitrogen  left  is  exactly  half  the  volume  of  the 
original  gas.  This,  however,  does  not  tell  us  anything  about 
the  volume  of  the  oxygen  in  the  compound  unless  we  know 
the  density  of  nitric  oxide.  When  the  gas  is  weighed,  it  is 
found  to  be  15  times  as  heavy  as  hydrogen.  Two  litres  of  it 
therefore  weigh  30  criths.  But  as  nitric  oxide  contains  half 
its  volume  of  nitrogen,  in  these  two  litres  there  must  be  one 
litre  of  nitrogen-  Now  one  litre  of  nitrogen  weighs  14  criths, 
therefore  we  get — 

2   litres   of    nitric   oxide,  weighing  30  criths 
contain  i  litre  of  nitrogen,  weighing  14     „ 

leaving  16  criths  as  the  weight  of 
the  oxygen ; 

but  1 6  criths  is  the  weight  of  i  litre  of  oxygen,  therefore  2 
litres  of  nitric  oxide  contain  i  litre  of  nitrogen  and  i  litre  of 


Nitrous  Oxide  201 

oxygen ;  or,  in  other  words,  2  volumes  of  nitric  oxide  contain 
i  volume  of  nitrogen  and  i  volume  of  oxygen. 

Nitrogen  Peroxide,  NO2. — This  is  the  red-brown  gas 
which  is  formed  when  nitric  oxide  comes  in  contact  with  air 

or  oxygen 

NO  +  O  =  NO2 

It  is  also  given  off  when  many  nitrates  are  heated. 

Experiment  173.— Heat  a  little  lead  nitrate  in  a  test-tube ;  notice 
that  the  tube  is  soon  filled  with  the  dark,  red-brown  gas.  We  have 
already  seen  that  this  gas  is  very  soluble  in  water,  therefore  we 
cannot  collect  it  over  water.  Attach  a  cork  and  delivery  tube  to 
the  test-tube,  and  see  if  any  gas  is  coming  off  besides  the  nitrogen 
peroxide,  by  collecting  it  over  water.  Note  that  a  colourless  gas 
is  collected.  Test  this  gas  with  a  glowing  splint  ;  the  gas  is 
oxygen.  Therefore,  oxygen  and  nitrogen  peroxide  are  both  pro- 
duced. The  equation  is — 

Pb(NO3)2  =  PbO  +  2NO2  +  O 

This  gas  plays  an  important  part  in  the  manufacture  of 
sulphuric  acid,  as  will  be  explained  later  on. 

Nitrous  Oxide,  N2O.— This  compound  is  easily  pre- 
pared by  simply  heating  the  salt  ammonium  nitrate. 

Experiment  174. — Place  some  crystals  of  ammonium  nitrate  in 
a  small  flask  fitted  with  a  cork  and  delivery  tube,  and  gently  heat 
over  a  small  flame.  The  crystals  quickly  melt  and  give  off  gas. 
Collect  the  gas  over  water  in  the  pneumatic  trough,  using  water  as 
warm  as  can  be  comfortably  borne  by  the  hands.  Collect  several 
jars  full.  The  equation  is  the  following — 

NH4NO3=  2H2O+  N2O. 

[Compare  this  reaction  with  that  by  which  nitrogen  was  pro- 
duced by  heating  ammonium  nitrite.] 

Properties  of  Nitrous  Oxide.— From  the  examples 
collected,  it  is  seen  to  be  a  colourless  gas.  It  has  a  faint, 
rather  pleasant  smell,  and  a  sweetish  taste.  The  gas  is 
moderately  soluble  in  cold  water,  but  less  so  in  warm  ;  there- 
fore there  is  not  so  much  loss  of  gas  if  warm  water  is  used  for 
its  collection.  When  nitrous  oxide  is  inhaled  for  a  short  time,  it 
causes  a  kind  of  intoxication,  often  accompanied  by  boisterous 


2O2  Oxides  of  Nitrogen. 

laughter.  On  this  account  it  is  often  called  laughing  gas.  If 
the  inhalation  is  continued  it  produces  a  state  of  insensibility 
to  pain,  and  for  this  reason  it  is  largely  used  by  dentists.1 

Experiment  175. — Bring  a  lighted  taper  into  a  jar  of  nitrous 
oxide.  Notice  that  the  gas  behaves  like  oxygen,  for  it  does  not 
burn,  but  causes  the  taper  to  burn  brightly.  Dip  a  glowing  splint 
of  wood  also  into  the  gas  ;  the  splint  is  rekindled  ;  so  that  from 
these  experiments  we  could  not  distinguish  this  gas  from  oxygen. 
Burn  a  piece  of  phosphorus  in  the  gas,  using  a  deflagrating  spoon 
as  in  Exp.  65.  The  phosphorus  burns  just  as  though  the  gas 
were  oxygen. 

How,  then,  can  we  distinguish  between  nitrous  oxide  and 
oxygen  ? 

Experiment  176. — Place  a  jar  of  oxygen  and  one  of  nitrous 
oxide  side  by  side.  Take  a  fragment  of  sulphur  on  a  deflagrating 
spoon,  or  on  a  bundle  of  asbestos  (see  Exp.  64),  light  one  corner 
of  it  in  a  gas  flame,  and  before  it  has  time  to  burn  up  at  all,  dip 
it  into  the  oxygen.  Notice  that  instantly  it  burns  up  vividly  and 
continues  burning.  Now  do  the  same  in  the  other  gas.  Notice 
what  happens ;  the  sulphur  is  extinguished.  Repeat  this  once  or 
twice  in  the  same  jar,  to  be  quite  sure  of  it.  Now  let  the  sulphur 
burn  up  well  before  plunging  it  into  the  gas,  and  notice  that  it 
continues  burning  when  thrust  into  the  nitrous  oxide. 

Nitrous  oxide,  then,  differs  from  oxygen  in  that  it  will 
extinguish  burning  sulphur  unless  the  sulphur  is  thoroughly 
hot  when  brought  into  it. 

Another  and  very  certain  way  of  distinguishing  between 
nitrous  oxide  and  oxygen,  is  to  mix  each  of  them  with  nitric 
oxide.  We  have  already  seen  what  happens  when  nitric  oxide 
is  mixed  with  oxygen  (Exp.  171).  But  when  nitric  oxide  is 
mixed  with  nitrous  oxide  there  is  no  production  of  brown 
fumes. 

Nitrous  oxide  and  oxygen  can  also  be  distinguished  by 
mixing  each  of  them  wi'th  hydrogen,  and  exploding  them.  For 
instance,  if  a  mixture  of  equal  volumes  of  oxygen  and  hydrogen 
is  exploded,  we  know  from  what  we  have  learnt  of  these  gases, 

1  Students  are  advised  not  to  experiment  upon  themselves,  or  each 
other,  by  inhaling  the  gas. 


Nitrous  Acid  and  Nitrites.  203 

that  the  hydrogen  will  combine  with  half  the  oxygen  to  form 
water,  and  that  half  the  original  volume  of  oxygen  will  remain 
over.  But  when  a  mixture  consisting  of  equal  volumes  of 
nitrous  oxide  and  hydrogen  is  exploded,  the  volume  of  gas  left 
is  the  same  as  the  original  volume  of  nitrous  oxide,  and  this  gas 
is  found  to  be  nitrogen. 

This  last  experiment  also  teaches  us  the  composition  of 
nitrous  oxide.  The  volume  of  nitrogen  present  is  equal  to  the 
volume  of  the  nitrous  oxide,  and  the  oxygen  present  was  just 
enough  to  unite  with  a  quantity  of  hydrogen  equal  in  volume 
to  the  nitrous  oxide.  We  know  that  this  amount  of  oxygen  is 
exactly  half  the  volume  of  hydrogen,  therefore  half  the  volume 
of  the  nitrous  oxide.  In  other  words,  two  volumes  of  nitrous 
oxide  contain  two  volumes  of  nitrogen  and  one  volume  of 
oxygen. 

Nitrous  Acid  and  Nitrites. — Nitrous  acid  has  never 
been  obtained  in  a  pure  state,  and  even  very  dilute  solutions 
of  it  quickly  decompose. 

The  formula  for  the  acid  is  HNOa }  it  has  one  atom  less 
oxygen  than  nitric  acid,  and  it  can  readily  take  up  oxygen  from 
compounds  which  are  rich  in  oxygen,  and  pass  into  nitric  acid. 

Experiment  177. — Dissolve  a  few  particles  of  sodium  nitrite  in 
half  a  test-tube  of  water,  and  add  to  it  a  few  drops  of  sulphuric 
acid.  This  decomposes  the  sodium  nitrite,  forming  sodium  sulphate 
and  nitrous  acid,  which  can  exist  for  a  short  time  in  the  dilute 
solution.  Now  add  to  this  some  solution  of  potassium  perman- 
ganate (a  salt  very  rich  in  oxygen),  and  notice  how  the  purple 
colour  of  the  permanganate  is  instantly  destroyed,  owing  to  the 
nitrous  acid  depriving  it  of  some  of  its  oxygen. 

On  the  other  hand,  nitrous  acid  will  give  up  some  of  its  own 
oxygen  to  many  substances  which  are  ready  to  unite  with  oxygen. 

The  salts  of  this  acid  are  called  nitrites,  as  sodium  nitrite, 
potassium  nitrite,  etc.  Potassium  nitrite  can  easily  be  made  by 
gently  melting  potassium  nitrate,  when  the  nitrate  parts  with  some 
oxygen  and  leaves  the  nitrite — 

KNO,  =  KNO2  +  O. 

We  can  easily  tell  a  nitrite  from  a  nitrate  by  adding  dilute 
sulphuric  acid  to  each.  The  nitrite  at  once  evolves  brown  fumes, 
the  nitrate  does  not. 


2O4  Oxides  of  Nitrogen. 

EPITOME. 
Five  oxides  of  nitrogen  are  known. 

1.  Nitrons  oxide  (laughing  gas),  N2O,  is  obtained  by  heating 
ammonium  nitrate. 

It  is  a  colourless,  slightly  sweet  tasting  gas.  Does  not  burn, 
but  supports  combustion  almost  as  well  as  oxygen.  Distinguished 
from  oxygen  by  (a)  much  greater  solubility  in  water,  (<$)  gives  no 
red  fumes  when  mixed  with  nitric  oxide. 

Nitrous  oxide  is  used  by  dentists  to  produce  insensibility  to 
pain.  It  is  called  laughing  gas,  because  when  a  little  of  it  mixed 
with  air  is  inhaled  it  produces  hysterical  laughter. 

2.  Nitric  oxide,   NO,  is  prepared   by  acting   on  copper  with 
nitric  acid.      Other  oxides  of  nitrogen  are  produced  at  the  same 
time,  although  at  a  certain  stage  of  the  action  the  bulk  of  the  gas 
given  off  is  nitric  oxide.     It  is  a  colourless  gas  which  instantly 
combines  with  free  oxygen  to  form  nitrogen  peroxide.     Therefore, 
when  it  comes  into  the  air  it  at  once  forms  brown  fumes  of  the 
peroxide.     Nitric  oxide  extinguishes  a  taper,  because  the  flame  is 
not  hot  enough  to  decompose  the  gas.     Even  phosphorus,  unless 
strongly  burning,  is  extinguished  ;  but  if  burning  vigorously  when 
introduced  into  the  gas,  the  flame  is  hot  enough  to  decompose  nitric 
oxide  into  oxygen  and  nitrogen,  when,  of  course,  the  phosphorus 
will  burn. 

Nitric  oxide  is  distinguished  from  all  other  gases  by  forming 
red  fumes  in  contact  with  air  or  oxygen. 

3.  Nitrogen  trioxide,  N2O3,  does  not  exist  as  a  gas.     When  a 
mixture  of  nitric  oxide,  NO,  and  nitrogen  peroxide,  NO2,  in  equal 
volumes,  is  passed  through  a  strongly  cooled  tube,  a  blue  liquid 
condenses,  which  is  believed  to  be  nitrogen  trioxide. 

4.  Nitrogen  peroxide,  NO2.      This  gas  is  formed  by  heating 
lead  nitrate.     When  the  gas  is  passed  through  a  cooled  tube,  it 
condenses   to   a   yellowish   liquid,   which    is    condensed   nitrogen 
peroxide.     The  liquid  gives  off  the  familiar  reddish  brown  gas. 

When  the  gas  is  heated,  it  rapidly  darkens  in  colour.  This 
change  is  due  to  the  molecules  breaking  up  into  simpler  ones.  At 
low  temperatures  nitrogen  peroxide  has  the  composition  N2O4,  its 
density  being  46  ;  but  as  the  temperature  gradually  rises,  the  gas 
breaks  up  into  molecules  having  the  composition  NO2,  and  at  140° 
its  density  is  23.  Therefore  nitrogen  peroxide  has  two  formulas, 
N2O4  at  low  temperatures,  and  NO2  at  high  temperatures.  Nitrogen 
peroxide  dissolves  in  water,  therefore  can  only  be  collected  by 
displacement. 


Oxides  of  Nitrogen.  205 

5.  Nitrogen  pentoxide,  N^O^,  is  obtained  by  withdrawing  the 
elements  of  water  from  pure  nitric  acid  by  means  of  phosphorus 
pentoxide. 

2HNO3  -  H2O  =  N2O5. 

It  is  a  white  crystalline  substance  which  cannot  be  preserved. 
It  combines  with  water  with  great  eagerness,  forming  nitric  acid. 
Nitrous  acid,  HNO2,  is  not  known  in  a  pure  state.     Its  salts 
are  nitrites. 


CHAPTER   XXII. 

OZONE. 

WHEN  we  hear  of  "  Smith  alias  Sampson  "  getting  into  trouble, 
we  know  at  once  that  the  two  names  stand  for  the  same  person. 
Very  often  when  a  man  assumes  an  alias  he  alters  his  outward 
appearance  as  far  as  he  can  in  order  to  make  it  more  difficult 
to  recognize  him.  Perhaps  from  being  a  fair  and  beardless 
person,  he  now  appears  as  a  dark  man  with  a  black  beard ; 
having  a  scar  on  his  face,  and  perhaps  even  a  squint. 

Now  some  of  the  elements  are  capable  of  doing  something 
of  the  same  kind.  Under  certain  circumstances  they  are  able 
to  assume  an  alias t  as  it  were.  They  adopt  new  outward 
appearances  and  fresh  properties  so  different  from  what  they 
usually  have,  that  they  are  then  altogether  unlike  their  original 
selves,  and  are  quite  likely  to  be  mistaken  for  some  entirely 
different  element  Oxygen  is  one  of  the  elements  which  can 
do  this.  With  the  usual  properties  of  this  element  we  are  now 
quite  familiar;  but  when  it  assumes  its  alias,  we  find  first  that 
it  adopts  a  powerful  and  unpleasant  smell.  Also  it  develops 
a  habit  of  attacking  organic  substances  in  a  violent  manner, 
so  that  it  cannot  even  be  passed  through  a  piece  of  indiarubber 
pipe  without  instantly  destroying  it.  Also,  it  attacks  metals 
like  silver  and  mercury,  which  it  took  no  notice  of  before ;  and 
if  it  comes  into  contact  with  potassium  iodide  it  instantly 
seizes  it,  steals  the  potassium  from  it,  and  turns  the  iodine 
adrift. 

We  see,  therefore,  that  when  oxygen  adopts  these  new 
hab*its  and  properties,  it  seems  like  an  entirely  different  sub- 
stance, although  all  the  time  it  is  nothing  else  but  oxygen. 


Ozone.  207 

Phosphorus  is  another  element  which  can  do  the  same 
thing.  As  usually  seen,  phosphorus  is  a  wax-like  solid,  very 
poisonous,  easily  cut  with  a  knife,  melts  when  warmed  by  the 
hand,  and  takes  fire  so  easily  that  it  has  to  be  kept  beneath 
water,  and  care  has  to  be  taken  in  handling  it.  When  it 
assumes  its  alias,  however,  it  becomes  a  dark  red  substance, 
which  looks  like  chocolate;  it  is  no  longer  poisonous,  does 
not  melt,  and  requires  to  be  strongly  heated  to  make  it  burn. 
All  the  time  it  is  phosphorus,  and  nothing  else  but  phosphorus, 
but  yet  so  unlike  the  element  in  its  ordinary  state. 

The  word  that  is  used  for  this  curious  property  is  allotropy 
(meaning  "  other  form ").  We  say  that  the  element  when 
appearing  in  its  alias,  or  its  more  unusual  character,  is  an 
allotropic  modification,  or  more  shortly,  an  allotrope  of  that 
element.  Thus,  when  phosphorus  appears  as  the  red  substance, 
we  call  this  the  allotropic  modification  of  phosphorus,  and 
when  oxygen  is  in  the  condition  in  which  it  shows  such  active 
properties,  we  speak  of  it  as  allotropic  oxygen,  or  ozone 
(ozone  means  *'  a  smell  "). 

We  can  make  oxygen  assume  this  allotropic  state  in  several 
ways.  The  best  method,  and  the  one  which  gives  the  largest 
amount  of  ozone,  is  to  expose  oxygen  to  a  particular  kind  of 
electrification,  known  as  the  silent  discharge. 

Experiment  178. — Take  a  piece  of  narrow  glass  tube  about  30 
centimetres  (12  inches)  long  and  wind  a  spiral  of  thin  copper  wire 
round  the  entire  length  outside  :  about  three  turns  of  wire  to  each 
centimetre. 

Attach  a  piece  of  indiarubber  tube  to  one  end,  so  as  to  conduct 
a  stream  of  oxygen  through.  Pierce  a  hole  with  a  pin  through  the 


FIG.  90. 

rubber  just  beyond  the  end  of  the  glass  tube,  and  push  a  thin  copper 
wire  through  it,  so  that  the  wire  reaches  nearly  to  the  other  end  of 
the  glass  tube,  as  shown  in  Fig.  90. 

The  inside  and  outside  wires  are  then  connected  to  a  Ruhm- 
korfs  coil.     When  the  coil  is  set  in  action,  there  will  be  no  visible 


208  Ozone. 

spark  jumping  across  between  the  wires,  but  a  vast  number  of  tiny 
sparks  passing  from  the  whole  length  of  one  wire  to  the  other.  If 
now  a  slow  stream  of  oxygen  be  passed  through  the  tube,  it  has  to 
"  run  the  gauntlet "  of  this  host  of  little  sparks,  and  this  causes  a 
part  of  the  oxygen  to  change  itself  into  ozone.  Why  it  does  this, 
nobody  knows. 

Test  for  Ozone. — We  test  for  ozone  by  making  use  of  the 
property  it  has  of  setting  free  the  iodine  contained  in  potassium 
iodide ;  and  in  order  that  we  may  see  when  the  iodine  is  so 
liberated,  we  have  starch  present.  The  instant  the  iodine  is 
set  free  it  unites  with  the  starch  and  produces  the  blue  com- 
pound described  on  page  134. 

Experiment  179. — Make  a  little  thin  starch  paste  (see  p.  134), 
and  add  to  it  a  little  potassium  iodide  solution.  Cut  strips  of  white 
paper  and  dip  them  into  this  mixture,  and  hang  them  up  to  dry. 
These  are  called  ozone  test  papers. 

Moisten  such  a  test  paper  and  hold  it  at  the  end  of  the  tube, 
where  ozone  is  being  formed  in  the  last  experiment.  Notice  that 
the  paper  is  instantly  turned  blue.  This  means  that  ozone  has  de- 
composed some  of  the  potassium  iodide,  and  that  the  free  iodine 
has  then  united  with  the  starch.  Notice  the  curious  smell  of  the 
ozone  which  passes  out  of  the  tube. 

Ozone  is  quickly  transformed  back  into  ordinary  oxygen  by- 
being  heated. 

Experiment  180. — Attach  to  the  end  of  the  "ozone  tube"  a 
short  straight  tube,  by  means  of  a  rather  wider  tube  and  two  corks, 


FIG.  91. 

as  shown  in  Fig.  91.     (The  usual  indiarubber  joint  cannot  be  used, 
for  the  reason  stated  above.)    Pass  the  ozonized  oxygen  through, 


Ozone.  209 

and  at  the  same  time  heat  the  tube.  The  ozone  will  be  changed 
into  ordinary  oxygen,  and  if  the  gas  passing  out  is  tested  with  a 
piece  of  ozone  test  paper,  there  will  be  no  blue  colour  produced. 

Chemists  now  know  that  the  change  which  oxygen  undergoes 
in  passing  from  oxygen  into  ozone,  is  that  the  molecule  of 
oxygen  takes  up  another  atom  of  oxygen.  Molecules  of  oxygen 
consist  of  two  atoms,  while  molecules  of  ozone  are  composed 
of  three  atoms  of  the  same  element.  Therefore,  when  oxygen 
changes  into  ozone  there  is  a  contraction  in  the  volume,  and  vice 
versA,  when  ozone  passes  back  into  ordinary  oxygen  there  is  an 
expansion  in  the  volume.  Ozone  is  on  this  account  sometimes 
spoken  of  as  "  condensed  oxygen."  This  is  quite  true  in  one 
sense,  because  the  molecule  of  ozone  is  more  condensed  than 
the  molecule  of  oxygen  ;  but  it  must  be  remembered  that  this 
is  quite  a  different  thing  from  condensing  or  compressing 
oxygen.  We  cannot  condense  oxygen  into  ozone  merely  by 
compressing  the  former  so  as  to  reduce  the  volume.  Ozone  is 
present  in  small  quantities  in  country  air.  It  is  produced  by 
lightning  discharges. 

It  is  also  formed  in  small  quantities  when  certain  substances 
combine  with  oxygen  without  actually  burning.  Thus,  if  a  piece 
of  phosphorus  is  placed  on  the  table,  it  is  seen  to  "  smoke." 
It  is  really  oxidizing  rather  quickly,  but  is  not  actually  burning. 
Now  the  phosphorus,  in  thus  combining  with  atmospheric 
oxygen,  causes  a  little  of  the  oxygen  to  go  into  the  allotropic 
form.  How  and  why  it  does  so,  is  not  known. 

Experiment  181. — Take  a  short  stick  of  phosphorus  (if  it  is 
coated  over  with  a  white  or  greyish  film,  scrape  it  clean,  underneath 
water]  and  place  it  in  a  good  large  bottle,  having  a  small  layer  of 
water  on  the  bottom— just  enough  to  half  bury  the  phosphorus. 
Cover  the  bottle  with  a  piece  of  cardboard  and  leave  it  for  about  ten 
minutes.  Then  dip  into  the  bottle  a  strip  of  moist  "  ozone  paper," 
and  note  that  it  shows  the  presence  of  ozone. 


CHAPTER  XXIII. 

CARBON. 

Occurrence. — Carbon  is  an  example  of  an  element  which 
can  assume  three  allotropic  forms. 

In  the  first  it  appears  as  a  soft  dull-black  solid,  which  has 
no  crystalline  shape.1  Charcoal  is  the  most  familiar  example 
of  this  form  of  carbon. 

The  second  variety  is  also  soft  and  black,  but  is  bright  and 
shiny,  almost  like  steel,  and  has  a  crystalline  form.  Graphite, 
or  "  black-lead,"  is  the  name  of  this  variety  of  carbon. 

The  third  modification  is  extremely  hard ;  sometimes  per- 
fectly colourless  and  transparent,  and  highly  crystalline.  The 
name  of  this  allotropic  form  of  carbon  is  Diamond. 

Each  of  these  three  allotropes  of  carbon  is  found  in  nature. 

Carbon  also  occurs  in  nature  in  a  state  of  chemical  com- 
bination with  other  elements.  For  example,  one  of  its  com- 
pounds with  oxygen,  namely,  the  gas  carbon  dioxide,  is  present 
in  the  air,  and  is  sent  out  in  large  volumes  from  rents  in  the 
rocks  in  volcanic  districts.  Again,  one  of  its  compounds  with 
hydrogen,  the  gas  known  as  marsh  gas,  or  fire-damp,  is  found 
in  large  quantities  in  coal  mines,  and  is  given  out  by  rotting 
vegetable  matter  in  marshy  places  (hence  its  name,  "  marsh 
gas"). 

Carbon  is  also  a  constituent  of  all  animal  and  vegetable 
matters ;  therefore  meat,  wool,  bread,  sugar,  alcohol,  wood, 
all  contain  carbon,  and  in  many  cases  the  carbon  is  associated 
simply  with  hydrogen  and  oxygen. 

1  Substances  which  show  no  crystalline  form  are  called  amorphous 
bodies — that  is,  without  form. 


Carbon. 


211 


Carbon  is  also  present  as  a  constituent  of  a  large  number 
of  the  common  minerals  which  help  to  make  up  the  solid 
earth.  Thus,  chalk,  marble,  limestone,  dolomite,  are  all  extremely 
common  and  abundant  minerals,  and  they  all  contain  carbon, 
associated  with  oxygen  and  lime,  as  carbonate  of  lime. 

How  to  obtain  Carbon  from  its  Compounds. — When 
a  piece  of  meat  or  bread  or  wood  is  partly  burnt,  we  say  that 
it  is  charred.  This  means  that  some  of  the  hydrogen  and 
oxygen  have  been  expelled  as  water,  and  a  portion  of  the 
carbon  has  been  set  free. 

Experiment  182. — Place  a  little  dry  powdered  starch  in  a  hard 
glass  tube,  the  end  of  which  has  been  sealed  up  (see  p.  37),  then 
draw  out  the  tube  and  bend 
it  as  shown  in  Fig.  92.  This 
constitutes  a  small  retort. 
Now  heat  the  starch,  and  no- 
tice that  it  chars  or  blackens, 
and  at  the  same  time  a  liquid 
is  expelled.  This  condenses 
in  the  drawn-out  part  of  the 
tube  and  can  be  collected  in 
a  small  test-tube.  This  liquid 
is  chiefly  water,  resulting  from 
the  decomposition  of  the 
starch ;  and  the  blackened 
mass  contains  free  carbon.  To 
identify  the  liquid  as  water, 
it  will  be  sufficient  here  to  FlG>  92< 

drop  upon  it  a  small  fragment 

of  potassium.      If  the  metal  takes  fire  (as  in  Exp.  44)  we   may 
safely  conclude  that  the  liquid  is  water. 

The  process  here  illustrated  is  that  of  destructive  distillation 
(see  p.  179),  and  whenever  compounds  containing  carbon  are 
submitted  to  this  treatment,  carbon  is  set  free  in  a  more  or 
less  pure  state,  depending  on  circumstances. 

Another  way  by  which  we  can  get  carbon  out  of  some  of 
its  compounds,  is  by  withdrawing  the  other  elements  with  which 
it  is  associated,  not  by  fire,  but  by  the  use  of  some  chemical 
reagent. 


212  Carbon 

For  example,  sugar  is  a  compound  of  carbon  with  oxygen 
and  hydrogen,  and  we  can  throw  the  carbon  out  of  combina- 
tion by  acting  on  the  sugar  with  sulphuric  acid. 

Experiment  183. — Roughly  weigh  out  12  grams  of  loaf  sugar  ; 
place  it  in  a  good-sized  beaker  and  pour  over  it  10  cc.  of  warm 
water.  This  will  dissolve  the  sugar  in  a  little  while,  especially  if 
gently  warmed,  and  give  a  strong  syrup.  Now  pour  into  this,  all 
at  once,  12  cc.  of  strong  sulphuric  acid  ;  the  whole  mixture  at  once 
froths  right  up.  The  acid  takes  from  the  sugar  the  hydrogen  and 
oxygen,  and  leaves  the  carbon,  which  will  appear  as  a  black  spongy 
mass. 

Alcohol,  or  "spirits  of  wine,"  is  another  compound  of 
carbon  with  hydrogen  and  oxygen,  but  it  has  a  great  deal 
more  hydrogen  in  proportion  to  carbon  than  sugar  has.  Let 
us  try  and  get  the  carbon  out  of  some  alcohol. 

Experiment  184. — Pour  15  cc.  of  strong  sulphuric  acid  into  a 
little  flask,  and  add  5  cc.  of  water.  Cool  it  by  dipping  the  flask  into 
cold  water.  Then  add  5  cc.  of  pure  alcohol.  [If  pure  alcohol  is 
not  available,  use  methylated  spirit.]  Provide  a  cork  and  delivery- 
tube  to  the  flask,  and  arrange  to  collect  gas  in  the  water  trough. 
Gently  heat  the  flask,  gas  soon  begins  to  come  off.  Collect  first  a 
small  jar  full,  and  allow  the  rest  to  collect  in  a  tall  cylinder  until 
it  is  about  one-third  filled  with  the  gas. 

[If  pure  alcohol  has  been  used,  there  will  be  very  little  blackening 
of  the  mixture  in  the  flask,  but  methylated  spirit  contains  other 
carbon  compounds  which  very  quickly  char,  or  carbonize  when 
treated  in  this  way.] 

The  gas  we  have  collected  is  called  ethylene.  It  contains 
some  of  the  carbon  from  the  alcohol,  combined  with  some  of 
the  hydrogen ;  so  that  by  this  experiment  we  have  not  yet  got 
free  carbon,  but  have  only  expelled  it  from  the  alcohol  still 
combined  with  hydrogen.  The  gas  is,  therefore,  a  hydro- 
carbon, as  it  is  called ;  that  is,  simply  a  compound  of  carbon 
and  hydrogen.  Hydro-carbons,  as  a  rule,  burn  with  a  flame 
which  gives  a  good  light. 

Experiment  185. — Remove  the  small  jar  of  gas  and  bring  a 
lighted  taper  to  it.  Notice  the  kind  of  flame  it  burns  with.  Con- 
trast this  flame  with  that  of  burning  hydrogen. 


Diamond.  2 1 3 

We  have  learnt  that  chlorine  has  the  power  of  taking 
hydrogen  away  from  compounds  of  carbon  and  hydrogen,  for 
in  Exp.  1 1 6,  when  turpentine  (which  is  also  a  hydro-carbon) 
was  brought  into  chlorine,  carbon  was  set  free.  Let  us,  there- 
fore, try  and  get  the  carbon  out  of  the  gas  we  have  obtained 
from  alcohol,  by  acting  on  it  with  chlorine. 

Experiment  1 86.  — Collect  a  quantity  of  chlorine  in  the  long 
cylinder  containing  the  gas  obtained  in  Exp.  184,  until  the  cylinder 
is  nearly  but  not  quite  full.  Slip  a  glass  plate  over  the  mouth  of 
the  cylinder  and  shake  up  the  remaining  water  in  it  so  as  to  mix 
the  gases  as  well  as  possible.  Then  apply  a  lighted  taper.  The 
mixture  burns  with  a  curious  flame,  producing  a  dense  black  smoke, 
and  a  black  deposit  of  carbon  all  down  the  sides  of  the  vessel. 
The  chlorine  has  combined  with  the  hydrogen  and  set  the  carbon 
free.  This  carbon  came,  therefore,  originally  out  of  the  alcohol. 

In  these  experiments  it  is  to  be  noted  that  the  carbon  is 
always  obtained  in  the  first  allotropic  modification,  that  is,  as 
a  soft  black  non-crystalline  substance,  like  charcoal  or  soot. 
It  is  very  much  more  difficult  to  obtain  carbon  in  the  second 
form,  as  graphite ;  and  to  make  it  pass  into  the  third  or 
diamond  variety  is  a  task  which  has  baffled  almost  every 
attempt. 

Diamond. — This  form  of  carbon,  although  the  most  valuable 
in  one  sense,  is  by  far  the  most  useless  so  far  as  the  chemist  is 
concerned.  It  is  found  in  gravel  deposits  in  India,  Africa,  and 
Australia.  When  found,  diamonds  are  not  at  all  like  the  gems. 
They  more  often  look  like  common  little  rough  stones,  scarcely 
transparent,  and  with  little  appearance  of  being  crystals.  To 
obtain  them  as  usually  seen  in  the  gem,  they  are  ground  or 
u  cut"  to  the  desired  shape,  so  as  to  " sparkle"  in  the  light. 

Diamond  is  the  hardest  known  substance,  and  will  scratch 
all  other  stones.  It  is,  therefore,  employed  for  cutting  glass, 
and  for  giving  a  hard  edge  to  drills  and  rock-borers.  Some 
diamonds  are  brownish,  and  even  black.  These  are  valueless 
for  gems,  and  are  used  for  drilling  and  also  for  grinding  and 
polishing  the  clear  ones.  When  strongly  heated  in  oxygen, 
the  diamond  first  turns  black  and  then  burns,  giving  carbon 


2 1 4  Carbon. 

dioxide  as  the  only  product  This  proves  that  it  is  pure 
carbon. 

Rock  crystal  or  quartz  is  sometimes  cut  to  look  like 
diamonds.  It  is  easy  to  distinguish  between  them  by  the 
fact  that  diamond  burns  in  oxygen  and  gives  carbon  dioxide, 
whereas  quartz  does  not  burn  at  all.  Quartz  is  not  nearly  so 
hard  as  diamond. 

Graphite. — This  form  of  carbon  is  not  nearly  so  rare 
as  diamond.  It  is  found  in  great  quantity  in  California.  It 
can  be  made  by  dissolving  charcoal  in  melted  iron  (just  as 
salt  is  dissolved  in  water)  and  then  allowing  the  iron  to  cool 
As  it  cools  carbon  crystallizes  out,  and  deposits  in  the  form 
of  graphite.  Long  before  it  was  known  that  graphite  was 
simply  carbon,  it  was  employed  for  writing  purposes.  It 
is  so  soft  that  when  drawn  across  paper  it  leaves  a  black 
mark.  Hence  it  got  the  name  plumbago,  or  "  the  writing  lead," 
and  it  is  also  known  under  the  common  name  of  black-lead. 
It  must  be  remembered,  however,  that  these  names  were 
given  to  it  simply  because  it  looked  rather  likz  lead.  Graphite 
is  used  for  making  "  lead  "  pencils,  and  for  "  black-leading  " 
iron  work.  When  strongly  heated  in  oxygen,  graphite,  like 
diamond,  burns,  and  gives  the  same  oxide  of  carbon. 

Amorphous  Carbon. — This  includes  common  wood- 
charcoal,  coke,  gas-carbon,  lampblack,  soot,  animal- charcoal 
or  boneblack.  All  these  substances  consist  of  non-crystalline 
carbon,  more  or  less  impure. 

Charcoal  is  made  by  subjecting  billets  of  wood  to  the 
process  of  destructive  distillation  (just  as  the  starch  was 
treated  in  Exp.  182)  in  ovens  or  retorts,  so  as  to  collect  the 
volatile  products  as  well  as  to  secure  the  charcoal :  or  else 
it  is  made  by  piling  the  wood  into  stacks,  and  setting  fire 
to  the  interior ;  the  outside  being  so  covered  over  with  turf 
or  earth  as  to  prevent  much  air  getting  to  the  smouldering 
heap.  By  this  method  everything  is  lost  except  the  charcoal. 
Care  has  to  be  taken  to  regulate  the  supply  of  air,  for  if 
too  much  gets  into  the  heap  the  charcoal  begins  to  burn  away 
and  is  also  lost. 

Charcoal  floats  when  thrown  on  water,  but  it  is  not  really 


Charcoal. 


215 


lighter  than  water ;  it  only  floats  because  its  pores  are  full  of 
air,'  and  this  buoys  it  up. 

Experiment  187. — Tie  a  little  weight  to  a  piece  of  charcoal  so 
as  to  sink  it,  and  throw  it  into  some  water  in  a  test-tube,  and  then 
heat  the  water.  Notice  bubbles  of  air  coming  out  of  the  charcoal. 
After  a  few  minutes'  boiling,  allow  the  water  and  charcoal  to  cool, 
and  it  will  be  found  that  it  will  no  longer  float  when  the  weight  is 
taken  off,  but  at  once  sinks  in  the  water. 

Charcoal,  like  both  the  other  forms  of  carbon,  burns  in 
oxygen  (see  Exp.  59),  and  produces  exactly  the  same  oxide  of 
carbon.  Charcoal  is  very  much  easier  to  burn,  however,  than 
either  diamond  or  graphite,  and  can  be  used  to  make  an 


FIG.  93. 

ordinary  fire  with.    It  would  be  quite  impossible  to  light  a  fire 
with  graphite  or  diamonds. 

Charcoal  has  a  wonderful  power  of  absorbing  gases,  and 
is  on  this  account  employed  to  arrest  the  bad  smelling  gases 
arising  from  drains  and  other  places. 

Experiment  188. — Fit  a  large  bottle  with  a  cork  and  two  glass 
tubes,  as  shown  in  Fig.  93,  B.  Break  up  some  charcoal  into  little 
pieces  and  after  heating  them  for  a  few  minutes  in  a  metal  dish 
or  tray,  fill  a  piece  of  combustion  tube  with  the  charcoal  and  fit  a 
cork  and  exit  tube  into  each  end. 

Now  place  a  few  small  particles  of  ferrous  sulphide  in  a  test- 
tube,  and  add  a  little  dilute  sulphuric  acid.  Remove  the  cork 


216  Carbon. 

from  the  bottle  B,  and,  by  means  of  a  piece  of  string,  lower  the 
test-tube  for  a  few  seconds  into  the  bottle  ;  then  draw  it  out  and 
replace  the  cork.  Attach  the  tube  W  to  a  water  tap,  and  place  a 
pinch-cock  on  G.  In  this  way  a  small  quantity  of  a  gas,  whose 
offensive  smell  will  have  been  noticed,  has  been  put  into  the  air 
in  the  bottle  B. 

This  gas  is  called  sulphuretted  hydrogen,  and  although  its 
smell  is  quite  enough  to  recognize  it  by,  we  can  use  a  more 
convenient  test. 

Dip  a  piece  of  paper  into  a  solution  of  lead  acetate  (sometimes 
called  sugar  of  lead},  and  let  a  little  of  the  air  in  the  bottle  blow 
against  it,  by  allowing  water  to  enter  through  tube  W,  and  opening 
the  pinch-cock  on  tube  G. 

Notice  that  the  paper  is  stained  black  by  the  gas. 

Now  connect  G  to  one  end  of  the  tube  containing  the  charcoal, 
and  gently  open  the  pinch-cock  so  as  to  drive  the  air  in  the  bottle 
slowly  through  the  charcoal.  Test  the  air  as  it  passes  out  of  the 
tube  with  another  piece  of  paper  moistened  with  acetate  of  lead. 
It  no  longer  is  blackened  :  neither  will  any  smell  of  sulphuretted 
hydrogen  be  detected.  The  charcoal  has  absorbed  all  the  bad 
smelling  gas. 

No  variety  of  wood  charcoal  is  quite  pure  carbon.  For 
instance,  if  we  burn  a  piece  of  charcoal  there  is  always  a 
certain  amount  of  white  ash  left.  This  of  course  is  not 
carbon ;  but  besides  this,  there  is  always  a  certain  amount  of 
some  compounds  of  carbon  with  hydrogen. 

Coke  is  made  by  treating  coal  very  much  as  wood  is 
treated  in  making  charcoal.  Coke  stands  very  much  in  the 
same  relation  to  coal  as  charcoal  does  to  wood.  Large 
quantities  of  it  are  produced  in  the  manufacture  of  ordinary 
coal  gas,  when  coal  is  heated  in  large  fireclay  retorts.  Coke 
is  often  manufactured  specially  for  the  coke  itself,  either  by 
burning  coal  in  great  stacks  (like  the  charcoal  stacks),  when 
all  the  gases  and  liquids  produced  at  the  same  time  are  lost ; 
or  in  special  ovens  or  kilns,  when  these  are  caught.  Coke, 
like  charcoal,  contains  a  certain  amount  of  mineral  matter, 
which  is  left  as  ash  when  the  coke  is  burnt;  and  it  also 
contains  a  little  hydrogen.  It  is  much  harder  and  heavier 
than  charcoal,  and  not  so  easily  lighted. 


A  nimal-  Charcoal.  2 1 7 

Gas-Carbon  is  a  still  harder  sort  of  coke,  which  is  found 
lining  the  retorts  in  which  coal-gas  is  manufactured.  It  is 
even  harder  and  heavier  than  ordinary  coke,  and  conducts 
electricity  extremely  well.  This  is  the  form  of  carbon  which 
is  employed  for  making  the  carbon  rods  used  in  electric  lights. 

Lampblack  is  the  soot  obtained  by  burning  substances 
like  petroleum  or  tar,  which  give  off  a  large  quantity  of  smoke. 
It  is  chiefly  used  for  printers'  ink,  and  for  black  paint. 

Animal- Charcoal  is  the  name  given  to  the  substance 
obtained  by  charring  bones  in  iron  retorts  (just  as  wood  or 
coal  is  treated).  It  is  a  very  impure  product,  containing  only 
about  10  parts  of  carbon  in  100  parts,  the  rest  being  the 
mineral  matters  of  the  bone  (chiefly  phosphate  of  lime). 
When  this  is  ground  up  fine  it  is  called  boneblack. 

This  substance  is  used  chiefly  for  removing  the  dirty 
colour  from  sugar  syrup,  in  making  white  sugar.  It  has  a 
great  power  of  absorbing  colouring  matter,  and  is  better  in 
this  respect  than  any  other  kind  of  charcoal. 

Experiment  189. — Half  fill  a  wide-mouthed  stoppered  bottle 
with  water,  which  has  been  tinted  with  either  magenta,  aniline  blue, 
indigo,  or  some  other  such  colouring  matter.  Add  to  this  a 
spoonful  of  fine  boneblack,  and  shake  it  well  for  a  moment  or 
two.  Then  filter  the  liquid,  and  note  that  what  passes  through  the 
filter  paper  is  quite  free  from  colour. 

All  the  varieties  of  amorphous  carbon  burn  in  oxygen 
and  give  carbon  dioxide,  the  same  product  as  graphite  or 
diamond  yield. 

Carbon  is  an  element  which  does  not  readily  enter  into 
chemical  union  with  other  elements.  Thus,  at  ordinary  tem- 
peratures, carbon  is  not  acted  on  by  hydrogen,  oxygen, 
chlorine,  or  nitrogen.  On  this  account,  wood  which  has 
been  charred  on  the  surface  is  not  so  soon  rotted  when  buried 
in  the  ground,  as  wood  which  has  not  been  so  protected. 
Hence  it  is  usual  to  char  the  end  of  stakes  or  posts  before 
putting  them  into  the  ground. 

Coal  is  the  product  of  a  process  of  natural  decomposition 
of  vegetable  matter  which  has  taken  long  ages  to  complete. 


218  Carbon. 

The  coal  we  burn  to-day  was  once  living  vegetation.  It  has 
long  been  buried,  owing  to  geological  changes  having  taken 
place,  and  has  been  subjected  to  pressure.  Coal  is  a  very 
impure  form  of  carbon,  and  contains  compounds  of  carbon 
with  hydrogen  and  oxygen.  Roughly  speaking,  the  proportion 
of  carbon  in  soft  or  bituminous  coal  is  about  80  in  100  parts; 
while  in  the  hard  or  anthracite  varieties  it  is  about  90  parts 
per  100. 

Although  carbon  is  not  acted  on  by  oxygen  at  ordinary 
temperatures,  it  unites  with  it  very  readily  when  heated.  Not 
only  will  it  combine  with  free  oxygen,  but  it  will  easily  take 
oxygen  away  from  certain  oxides.  On  this  account  carbon  is 
a  most  useful  element  to  the  metallurgist,  for  by  heating  oxides 
of  metals  with  carbon,  the  carbon  takes  the  oxygen  and  leaves 
the  metal  in  the  uncombined  state. 

Experiment  190. — Take  a  small  quantity  of  red-lead,  Pb3O4,  and 
mix  it  thoroughly  with  about  one  quarter  as  much  finely  powdered 
charcoal.  Put  the  mixture  in  a  small  porcelain  crucible  and 
strongly  heat  it.  After  a  short  time  throw  out  the  contents  of  the 
crucible,  and  there  will  be  a  globule  of  bright  metallic  lead. 

This  process  is  called  reduction.  We  say  that  the  lead 
has  been  reduced  by  the  charcoal.  This  simply  means  that 
the  oxygen  with  which  the  lead  was  combined  has  been  taken 
from  it  by  the  carbon,  and  the  lead,  therefore,  was  left  by 
itself. 

This  is  how  many  of  the  metals  are  obtained  from  their 
ores.  Iron  ores,  for  instance,  which  are  used  for  getting  iron 
from,  are  oxides  of  iron.  These  when  very  strongly  heated 
with  carbon  (usually  coal  or  coke)  are  deprived  of  their 
oxygen,  and  the  iron  is  set  free.  All  metals  cannot  be  reduced 
from  their  oxides  in  this  way. 

EPITOME. 

Carbon  is  an  element  known  in  three  allotropic  forms,  (i) 
Diamond.  (2)  Graphite.  (3)  Charcoal. 

Diamond  is  the  hardest  known  substance  ;  highly  crystalline, 
and  often  nearly  or  quite  colourless.  Used  as  a  gem,  and  for  cutting 
purposes. 


Carbon,  219 

Graphite  is  soft,  black,  shining,  and  crystalline.  Used  for  "  lead  " 
pencils,  for  "  black-leading." 

Charcoal  is  soft,  black,  dull,  and  non-crystalline.  Used  for  fuel 
and  for  making  gunpowder. 

Coke,  lampblack,  soot,  bone  black  or  animal-charcoal  are  all 
more  or  less  impure  forms  of  amorphous  carbon. 

All  the  varieties  of  carbon  burn  in  oxygen  and  give  the  same 
compound,  namely,  carbon  dioxide.  In  this  way  diamond  is  easily 
distinguished  from  quartz  (which  is  silicon  dioxide,  and  will  not 
burn)  or  other  imitations. 

Carbon  in  combination  occurs  in  all  organic  substances,  and 
when  these  are  charred,  or  heated  so  that  air  does  not  get  to  them 
to  cause  them  to  burn  away  altogether,  the  carbon  is  left  in  a  more 
or  less  pure  state.  Wood-charcoal,  animal-charcoal,  and  coke, 
are  obtained  by  heating  wood,  bones,  and  coal  in  this  manner. 

Charcoal  absorbs  gases  readily,  and  is  therefore  used  to  remove 
bad  smelling  gases  ;  wood-charcoal  does  it  best. 

Charcoal  also  absorbs  colouring  matter  from  liquids  which  are 
filtered  through  it  or  shaken  up  with  it.  Animal-charcoal  does 
this  best ;  it  is  therefore  used  for  filters. 

Coal  is  an  impure  form  of  carbon,  containing  a  number  of 
carbon  compounds. 

Charcoal,  coke,  and  coal  are  used  as  reducing  agents,  for 
taking  oxygen  away  from  oxides  of  certain  metals  like  iron,  copper, 
lead,  etc.  When  such  oxides  are  strongly  heated  with  carbon,  the 
metal  they  contain  is  reduced. 

This  is  the  principle  of  the  smelting  of  many  of  the  ores  of 
metals. 


CHAPTER  XXIV. 

CARBON  DIOXIDE,  CARBONIC  ACID,  CARBONATES. 

Carbon  Dioxide,  CO2. — When  a  piece  of  charcoal  is 
burnt  in  oxygen,  as  in  Exp.  63,  the  carbon  and  oxygen  combine 
and  form  the  compound  called  carbon  dioxide.  Let  us  repeat 
that  experiment  in  a  slightly  different  way,  so  as  to  collect  the 
compound  and  examine  it. 

Experiment  191. — Place  a  few  little  pieces  of  charcoal  in  a  short 
piece  of  combustion  tube,  through  which  a  stream  of  oxygen  from 


FIG.  94. 

a  gasholder  can  be  passed  ;  attach  a  cork  and  delivery  tube  to  one 
end,  and  arrange  to  collect  the  gas  over  water,  as  shown  in  Fig.  94. 
Now  heat  the  charcoal  by  means  of  a  Bunsen  flame  and  allow  it 
to  burn  in  the  stream  of  oxygen,  which  should  be  regulated  so  as 
just  to  keep  the  charcoal  burning  brightly.  Collect  two  jars  of 
the  gas. 

Into  one  jar  dip  a  lighted  taper.  Notice  that  the  flame  is 
instantly  put  out,  and  also  that  the  gas  does  not  burn.  In  these 
two  respects  it  is  like  nitrogen. 


Carbon  Dioxide. 


221 


Now  pour  into  another  jar  some  clear  lime  water,1  and  shake 
it  up.  Note  that  the  lime  water  at  once  becomes  milky. 

This  action  of  carbon  dioxide  on  lime  water  is  a  test  by 
which  we  can  distinguish  this  gas  from  all  others,  as  it  is  the 
only  gas  which  can  produce  this  effect.  The  equation  is — 


CaH2O2 

Lime  water. 


coc 


=      CaCO3      -f 

Calcium  carbonate. 


FIG.  95. 

Not  only  is  carbon  dioxide  produced  when  charcoal  is 
burnt  in  oxygen,  but  when  any  carbon  compound  is  burnt  in 
the  air. 

Experiment  192. — Set  fire  to  a  piece  of  paper,  and  drop  it  while 
burning  into  a  dry  bottle,  and  cover  the  mouth  with  a  piece  of  card 
or  glass  plate.  Then  pour  in  a  little  lime  water,  and  shake  it  up. 

Experiment  193. — Hold  a  dry  jar  over  a  candle  flame  for  a 
minute  or  two,  so  as  to  catch  some  of  the  invisible  products  of  its 
burning,  as  in  Fig.  95.  Then  add  lime  water,  and  shake  up.  Do 
the  same  with  a  spirit  lamp  flame,  and  with  an  ordinary  coal-gas 

1  To  make  lime  water,  place  a  handful  of  powdered  lime  in  a  big 
bottle,  and  fill  it  up  with  water,  cork  it  up,  and  after  shaking  for  a  minute, 
leave  it  to  settle.  Then  pour  off  the  clear  liquid  into  a  separate  bottle  for 
use.  The  other  bottle  can  be  filled  up  with  water  over  and  over  again, 


222  Oxides  of  Carbon. 

flame.     Notice   that   in   every  case  there  is  abundance  of  carbon 
dioxide  being  produced. 

Roughly  speaking,  every  ton  of  coal  that  is  burnt,  produces 
about  three  tons  of  carbon  dioxide,  all  of  which  escapes  into 
the  atmosphere. 

We  have  already  learnt  that  respiration  is  a  sort  of  com- 
bustion, let  us  see  if  carbon  dioxide  is  produced. 

Experiment  194. — Place  some  lime  water  in  a  large  test-tube, 
and  by  means  of  a  glass  tube  make  the  breath  bubble  through  the 
solution.  Notice  that  the  first  portions  of  breath  produce  hardly 
any  milkiness,  but  presently,  when  air  that  has  been  further  into  the 
lungs  is  blown  through,  the  liquid  quickly  shows  that  there  is 
carbon  dioxide.  (See  also  Exp.  75.) 

Carbon  dioxide  is  also  given  off  during  many  processes  of 
decay  and  fermentation.  Thus,  when  vegetable  matter  (such 
as  the  leaves  of  trees  which  fall  in  autumn)  gradually  rots  away, 
carbon  dioxide  is  given  off. 

Experiment  195. — Dissolve  a  handful  of  sugar  (common  brown 
moist  sugar  is  best)  in  some  tepid  water  in  a  good-sized  flask,  and 
add  some  yeast.  Fit  a  cork  and  delivery-tube  to  the  flask.  The 
sugar  soon  begins  to  ferment  and  gas  escapes.  Collect  some  of  the 
gas  at  the  water  trough  and  test  it  with  lime  water.  In  breweries 
enormous  quantities  of  carbon  dioxide  are  in  this  way  produced. 

Since  these  different  operations,  namely,  the  burning  of 
ordinary  fuel,  respiration  of  man  and  animals,  decay  and 
fermentation,  besides  others  which  are  constantly  going  on, 
all  result  in  sending  carbon  dioxide  into  the  air,  it  is  no 
wonder  that  this  compound  is  present  in  the  atmosphere. 
The  wonder  rather  seems  that  the  air  does  not  become  too 
impure  to  breathe.  But,  strange  to  say,  if  we  test  a  small 
sample  of  ordinary  air  by  means  of  lime  water,  we  find  that 
we  can  scarcely  even  detect  the  presence  of  this  gas  at  all. 

Experiment  196.— Take  a  bottle  (such  as  has  been  used  for 
testing  in  the  above  experiments)  which  contains  simply  ordinary 
air.  Pour  some  lime  water  in  and  shake  it  up.  No  turbidity  of 
the  lime  water  is  noticeable.  Of  course,  if  this  experiment  is  made 
in  a  small  room  where  a  number  of  persons  have  been  present  for 


Carbon  Dioxide.  223 

some  time,  and  several  gas-lamps  have  been  burning,  most  likely 
the  air  will  contain  enough  carbon  dioxide  to  cause  a  turbidity  in 
lime  water  when  a  small  sample  of  it  is  tested  in  this  way. 

Since  in  the  open  air  there  is,  after  all,  so  small  an  amount 
of  carbon  dioxide  (only  about  3  parts  in  10,000  parts  of  air,  by 
measure),  in  spite  of  the  enormous  quantities  which  are  every 
day  being  sent  into  the  atmosphere — think  of  the  millions  of 
people  in  London  alone,  and  the  hundreds  of  tons  of  coal 
daily  burnt ! — it  is  evident  that  there  must  be  some  natural 
process  at  work  which  constantly  removes  this  compound  from 
the  atmosphere.  Nature,  as  usual,  has  made  a  beautiful 
provision  for  preventing  the  accumulation  of  carbon  dioxide, 
and  her  secret  for  doing  this  has  long  been  found  out.  She 
has  endowed  the  green  parts  of  plants  with  the  power  of 
decomposing  carbon  dioxide,  by  the  aid  of  sunlight,  into  its 
two  component  elements,  carbon  and  oxygen.  Every  green 
leaf  and  every  blade  of  grass  is  a  tiny  chemical  laboratory, 
where  oxygen  is  being  made  from  carbon  dioxide,  and  is 
ever  being  returned  to  the  atmosphere  with  all  its  life-support- 
ing properties.  The  carbon  which  comes  out  of  the  carbon 
dioxide  is  utilized  by  the  plant.  Vegetation,  therefore, 
removes  dioxide  from  the  atmosphere,  and  at  the  same  time 
gives  back  the  oxygen  which  had  been  taken  from  it  by  the 
carbon. 

To  prepare  Carbon  Dioxide  for  experiments,  we  usually 
adopt  quite  a  different  method. 

Experiment  197. — Take  the  apparatus  shown  in  Fig.  39,  p.  48, 
and  put  into  the  bottle  a  quantity  of  broken-up  marble  and  some 
water.  Now  pour  a  little  strong  hydrochloric  acid  down  the  funnel, 
and  note  that  effervescence  at  once  begins.  After  a  minute  or  two, 
when  the  air  is  all  swept  out,  collect  the  gas  which  is  given  off. 

Marble  is  a  variety  of  calcium  carbonate,  and  the  change 
here  going  on  is  the  following — 

CaCO3      +      2HC1      =      CaCl2      +     H2O     +     CO* 

Calcium  carbonate.  Calcium  chloride. 

Marble  is  used  because  it  is  one  of  the  most  convenient  of 


224  Oxides  of  Carbon. 

all  the  carbonates ;  but  we  can  equally  well  obtain  the  gas  by 
acting  on  any  carbonate  with  almost  any  common  acid. 

Experiment  198. — Break  up  some  sodium  carbonate  (common 
"  washing  soda  ")  and  put  a  little  into  three  small  beakers.  Stand 
each  of  the  beakers  in  a  larger  beaker,  or  jar,  and  then  by  means  of 
a  pipette,  add  to  one  a  few  drops  of  dilute  sulphuric  acid  ;  to  the 
next,  dilute  hydrochloric  acid  ;  and  to  the  third,  dilute  nitric  acid. 
Cover  the  outer  jars  with  pieces  of  paper.  The  little  beakers  can 
now  be  lifted  out  with  tongs,  and  lime  water  poured  into  each  jar. 
In  each  case  the  lime  water  will  become  turbid. 

In  practice  we  cannot  use  sulphuric  acid  with  calcium 
carbonate,  unless  the  carbonate  is  first  powdered  up  very  fine 
and  made  into  a  paste  with  water. 

Experiment  199. —  Put  a  few  lumps  of  marble  in  a  test-tube  with 
some  water,  and  add  a  little  strong  sulphuric  acid.  Notice  at  first 
there  is  an  effervescence,  but  that  it  very  soon  leaves  off,  and  no 
more  gas  comes  off.  The  first  action,  when  it  is  effervescing, 
is  this — 

CaC03       +       H2SO4       =       CaS04      +       H2O       +       CO2. 

Calcium  sulphate. 

Calcium  sulphate  is  formed;  this  is  the  same  as  "plaster 
of  Paris,"  and  it  at  once  coats  over  the  lump  of  marble  and 
prevents  any  more  acid  getting  to  it. 

Carbon  dioxide  is  a  very  heavy  gas.  It  is  about  i£  times 
heavier  than  air,  so  that  it  is  quite  easy  to  collect  a  jar  of  it 
by  "  downward  displacement."  More  than  this,  the  gas  can 
even  be  poured  from  one  jar  to  another  like  a  liquid ;  and 
if  poured  into  a  vessel  suspended  on  a  balance,  it  will  weigh 
down  that  end  of  the  beam. 

Experiment  200. — Take  a  long  light  wooden  rod,  and  hang  from 
one  end  of  it  a  light  cardboard  box  (or  an  old  hat),  and  place  upon 
the  other  end  a  piece  of  bent  lead  rather  lighter  than  the  box. 
Balance  this  upon  the  edge  of  a  paper-knife  or  other  convenient 
metal  edge,  in  the  manner  shown  in  Fig.  96.  Then  bring  a  good 
large  jar  of  carbon  dioxide,  and  pour  it  into  the  box  or  hat,  and 
notice  that  it  weighs  it  quite  down. 

Carbon  dioxide  is  soluble  to  a  small  extent  in  water,  but 
not  to  such  a  degree  as  to  make  any  difference  when  we 


Carbon  Dioxide.  225 

collect  the  gas  at  the  water  trough.  All  natural  waters  contain 
a  little  carbon  dioxide  dissolved  in  them,  and  some  waters 
contain  quite  a  large  quantity.  For  instance,  Seltzer  water 
and  Apollinaris  water  contain  so  much  of  this  gas  dissolved 
in  them  that  they  actually  effervesce  or  "  sparkle,"  owing  to 
the  escape  of  the  carbon  dioxide.  Under  ordinary  conditions 
water  dissolves  about  its  own  volume  of  this  gas ;  but  under 
increased  pressure  it  can  take  up  more.  When  such  extra 
pressure  is  again  removed,  the  extra  gas  that  was  dissolved 
is  given  off.  This  is  illustrated  in  the  ordinary  aerated  waters 
used  for  drinking.  Gas  is  pumped  into  the  water  under  great 


FIG.  96. 

pressure  so  that  the  water  dissolves  a  large  quantity,  but  the 
moment  the  pressure  is  released  by  drawing  the  cork,  then 
the  gas  rapidly  escapes  with  the  familiar  effervescence. 

Experiment  201. — Colour  some  ordinary  water  with  a  few  drops 
of  litmus,  and  let  carbon  dioxide  bubble  through  it  for  a  few 
minutes  from  the  apparatus  used  in  Exp.  197.  Notice  that  the 
litmus  turns  red,  showing  the  presence  of  an  acid.  Note  that  the 
colour  is  not  quite  so  bright  red  as  when  a  drop  of  hydrochloric 
or  dilute  sulphuric  acid  is  added.  It  is  because  the  solution  of 
this  gas  in  water  is  acid  that  the  gas  is  sometimes  called  "  carbonic 
acid  gas." 

The  acid  so  obtained  is  a  very  feeole  acid,  and  easily 
decomposes.  It  is  called  carbonic  acid,  H2CO3.  If  it  is 
gently  heated  it  is  decomposed  again  into  water  and  carbon 
dioxide. 

H2CO3  =  H20  +  C02. 

Q 


226  Oxides  of  Carbon. 

Experiment  202. — Take  some  of  the  solution  used  in  the  last 
experiment  and  heat  it  in  a  test-tube.  Notice  that  the  reddish 
colour  of  the  litmus  soon  changes  back  again  to  the  original  blue, 
as  the  feeble  acid  is  being  decomposed. 

Carbonates. — Although  real  carbonic  acid  is  only  a  feeble 
acid,  and  so  easily  decomposed  that  it  cannot  be  obtained  by 
itself,  it  forms  important  salts  called  carbonates. 

Calcium  carbonate,  CaCO3,  occurs  as  marble,  limestone, 
chalk.  This  is  the  compound  that  is  formed  when  lime  water 
is  brought  into  carbon  dioxide. 

Experiment  203. — Pass  some  carbon  dioxide  from  the  gene- 
rating apparatus  (Exp.  197)  into  some  lime  water  until  there  is  a 
thick  milkiness.  Filter  the  liquid  through  a  small  filter.  The 
white  deposit  will  hardly  be  visible  on  the  paper,  but  we  can  test 
it  in  the  following  way.  Put  the  paper  into  a  small  beaker,  and 
pour  upon  it  a  few  drops  of  hydrochloric  acid.  Notice  effervescence. 
Put  a  little  lime  water  in  a  wide  test-tube,  and  pour  some  of  the 
gas  out  of  the  beaker  into  it  and  shake  it  up. 

The  equation  which  shows  what  takes  place  when  carbon 
dioxide  is  passed  into  lime  water  is — 

Ca(HO)2       +       CO2       =       CaCO3       4-       H2O. 

Lime  water  or  calcium  Calcium  carbonate 

hydroxide.  or  chalk. 

so  that  we  get  back  the  same  compound,  calcium  carbonate, 
which  was  used  to  make  the  carbon  dioxide  from. 

Experiinent  204. — Pass  carbon  dioxide  through  a  solution  of 
caustic  soda.  Notice  that  there  is  no  turbidity  ;  also  observe  that 
the  gas  is  eagerly  absorbed,  because  bubbles  hardly  rise  to  the 
top  of  the  liquid.  Continue  bubbling  the  gas  into  the  solution 
until  no  more  is  absorbed,  and  then  gently  evaporate  the  solution. 
Take  a  little  of  the  residue  and  add  hydrochloric  acid  to  it.  Observe 
the  effervescence.  Test  the  gas  with  lime  water. 

Take  another  portion  of  the  residue  and  add  water  to  it  ;  it 
dissolves  easily.  This  is  the  reason  why  there  was  no  precipitate 
formed  when  the  gas  was  passed  into  caustic  soda  ;  the  sodium 
carbonate  that  was  formed  is  soluble  in  water.  The  equation  in 
this  case  is — 

2NaHO       +       CO2       =       Na,CO3       +       HaO. 

Caustic  soda  or  Sodium  carbonate, 

sodium  hydroxide 


Carbonates. 


227 


There    are   two    sodium    carbonates.      One  is   common 
"washing  soda,"  and  the  other  is  usually  called  sodium  bi- 
carbonate.     The   difference  in  these  salts  is  that   the  first 
contains  twice  as  much  sodium  as  the  last ;  or,  in  other  words, 
the  last  contains  twice  as  much  carbonic  acid  in  proportion 
to  sodium  as  the  first.     Their  formulae  are — 
Na2CO3,  sodium  carbonate,  or  normal  sodium  carbonate. 
HNaCO3,  sodium  bi-carbonate,  or  hydrogen  sodium  carbonate. 

Experiment  205. — Dilute  some  lime  water  with  about  half  as 
much  distilled  water,  and  pass  carbon  dioxide  through  it.  As  in  the 
former  experiment  there  is  a  precipitate  of  calcium  carbonate. 


FIG.  97. 

But  continue  passing  the  gas,  stirring  up  the  liquid,  and  notice 
that  in  a  few  minutes  it  becomes  perfectly  clear  again.  The 
calcium  carbonate  has  dissolved  in  the  solution  of  carbon  dioxide. 
Take  a  little  of  this  clear  solution  and  boil  it.  It  again  becomes 
turbid,  because  the  carbon  dioxide  is  expelled,  and  the  calcium 
carbonate  cannot  remain  in  solution.  The  presence  of  calcium 
carbonate  dissolved  in  this  way  is  what  causes  the  temporary  hard- 
ness of  natural  waters  (see  pp.  82  and  269). 

Many  carbonates  part  with  carbon  dioxide  when  they  are 
heated,  leaving  an  oxide. 

Experiment  206. — Heat  a  little  magnesium  carbonate  in  a  test- 
tube,  and  allow  the  gas  to  flow  down  into  another  test-tube,  con- 
taining lime  water,  in  the  manner  indicated  in  Fig.  97, 
MgCO3  =  CO3  +  MgO. 


228  Oxides  of  Carbon. 

Calcium  carbonate  undergoes  the  same  change,  but  re- 
quires a  higher  temperature. 

The  process  of  lime  burning,  carried  on  in  lime  kilns, 
illustrates  this. 

Limestone  (that  is,  calcium  carbonate)  is  heated  in  the 
kiln,  when  carbon  dioxide  escapes  into  the  air,  and  lime  (that 
is,  calcium  oxide)  remains. 

CaCO3   =   CO,  +  CaO. 

Limestone.  Lime. 

The  "  setting  "  of  mortar  is  partly  due  to  the  absorption 
of  carbon  dioxide  out  of  the  air  by  the  lime  in  the  mortar. 
Old  mortar,  therefore,  contains  calcium  carbonate. 

Experiment  207. — Pour  a  little  hydrochloric  acid  upon  some 
fragments  of  old  mortar  •  notice  the  effervescence,  and  pass  the 
gas  into  lime  water. 

To  find  the  Weight  of  Carbon  Dioxide  in  Marble.— 

Experiment  208. — Select  a  small  thin  flask  with  a  wide  mouth, 
and  fit  it  with  a  cork  having  two  holes.  Through  one  hole  insert 
a  short  straight  glass  tube,  the  top  of  which 
can  be  closed  with  a  tiny  cork.  Into  the 
other  hole  fit  a  bent  glass  tube  with  a  small 
bulb  blown  on  it,  as  shown  in  Fig.  98.  Make 
a  hole  in  the  bottom  of  a  small  test-tube,  by 
heating  the  end  with  a  fine-pointed  blow- 
pipe flame,  and  blowing  down  the  test-tube. 
Then  cut  off  the  tube  about  3^  cm.  (\\ 
inches)  from  the  bottom,  and  "  border  "  the 
end  (p.  36)  so  as  to  obtain  the  little  appa- 
ratus shown  at  A,  Fig.  98.  Fasten  a  thin 
copper  wire  to  the  neck,  and  hang  the  little 
tube  inside  the  flask  in  the  manner  shown. 

\  *:*-" ~  -       Place  a  little  strong  sulphuric  acid  in  the 

bulb  tube  B,  so  as  to  just  fill  the  bend,  and 
close  the  open  end  with  a  little  cap,  C 
(p.  in).  Put  about  10  or  12  cc.  of  a  mixture  of  equal  parts  of 
hydrochloric  acid  and  water  in  the  flask,  but  do  not  let  it  touch  the 
little  hanging  tube.  Now  carefully  weigh  the  whole  apparatus. 
Then  remove  the  cork  and  drop  into  the  hanging  tube  one  or  two 
little  fragments  of  marble  (not  so  small  as  to  drop  through  the  hole 


Analysis  of  a  Carbonate.  229 

in  the  bottotri),  and  weigh  again.  The  increase  gives  the  weight  of 
marble,  which  should  be  from  i  to  2  grams. 

Remove  the  cap  C,  and  then  lower  the  tube  containing  the 
marble  so  that  it  dips  into  the  hydrochloric  acid,  and  replace  the 
little  cork  d. 

The  acid  dissolves  the  marble,  and  the  carbon  dioxide  which  is 
evolved  makes  its  escape  by  bubbling  through  the  sulphuric  acid 
in  B.  This  prevents  the  gas  from  carrying  away  any  vapour  of 
water  with  it. 

When  the  marble  has  all  dissolved,  the  apparatus  of  course  is 
full  of  carbon  dioxide,  and  as  this  is  much  heavier  than  air  it  must 
be  removed.  Take  out  cork  d,  and  by  means  of  a  short  piece  of 
rubber  tube  attached  to  the  bulb-tube,  slowly  suck  the  gas  out  of 
the  flask,  making  it  bubble  quite  slowly  through  the  sulphuric  acid 
in  the  tube.  When  the  gas  is  all  out,  and  the  apparatus  is  full  of 
air  as  it  was  at  first,  replace  the  cork  d,  and  cap  C,  and  weigh  again. 
The  loss  of  weight  represents  the  carbon  dioxide  which  has  passed 
out.  Two  experiments  should  be  made,  so  as  to  confirm  the  result. 

EXAMPLE.         Weight  of  marble  used  =  1*5  grams 
Weight  of  carbon  dioxide  =  O'6      „ 

as  1*5  :  100  ::  0*6  =  40 

therefore  this  sample  of  marble  contains  40  per  cent,  of  carbon 
dioxide. 

EPITOME. 

Carbon  dioxide  is  produced  when  all  carbon  compounds  burn. 
Coal,  gas,  candles,  oil,  wood,  etc.,  when  burnt  give  carbon  dioxide 
amongst  other  things.  The  breathing  of  animals  and  man  pro- 
duces carbon  dioxide,  and  it  is  given  off  during  processes  of  decay 
and  fermentation. 

It  is  also  formed  when  limestone  (that  is,  calcium  carbonate) 
is  heated,  as  in  the  process  of  lime-making  in  lime  kilns. 

Carbon  dioxide  is  prepared  by  acting  on  calcium  carbonate 
(marble,  chalk,  or  limestone)  with  hydrochloric  acid. 

Carbon  dioxide  is  a  colourless  gas,  slightly  soluble  in  water 
forming  a  feebly  acid  solution,  which  is  decomposed  again  when 
heated.  The  gas  does  not  burn,  and  puts  out  ordinary  flames. 
Carbon  dioxide  is  not  exactly  poisonous,  because  we  are  always 
breathing  small  quantities  of  it  in  the  air,  but  if  animals  are  placed 
in  the  gas  they  quickly  die  for  want  of  oxygen.  Even  a  moderate 

1  The  student  should  calculate  what  weight  of  carbon  dioxide  would 
have  been  obtained  if  the  marble  had  been  pure  calcium  carbonate. 


230  Oxides  of  Carbon. 

quantity  of  this  gas  in  the  air,  over  and  above  the  usual  amount,  is 
injurious  to  life. 

Carbon  dioxide  is  a  heavy  gas,  which  can  be  collected'  by  down- 
ward displacement,  and  can  be  poured  from  one  vessel  to  another. 

When  passed  into  lime  water  it  unites  with  the  lime  and  forms 
a  white  precipitate  of  chalk.  It  is  quickly  absorbed  by  either 
caustic  soda  or  potash,  forming  sodium  or  potassium  carbonates. 

The  test  for  a  carbonate  is  to  add  an  acid  to  it,  and  to  prove  that 
the  gas  which  is  given  off  is  carbon  dioxide  by  passing  it  through 
lime  water  and  obtaining  the  precipitate  of  chalk. 


CHAPTER   XXV, 

OXIDES  OF  CARBON — continued. 

Carbon  Monoxide,  CO. — It  has  been  shown  that  when 
carbon  burns  in  air  or  oxygen  under  ordinary  circumstances 
it  gives  carbon  dioxide,  but  under  particular  conditions  the 
other  oxide  is  formed.  For  instance,  if  we  fill  a  long  piece  of 
combustion  tube  with  fragments  of  charcoal,  and  heat  the  tube 
to  a  good  red  heat  in  a  furnace,  and  then  pass  a  moderate 
stream  of  oxygen  in  at  one  end,  instead  of  getting  carbon 
dioxide  coming  out  at  the  other,  we  should  find  that  it  was 
quite  a  different  gas ;  for  we  could  set  fire  to  it,  and  should  see 
it  burn  with  a  beautiful  blue  flame.  What  takes  place  is  this. 
As  the  oxygen  first  meets  the  hot  charcoal,  the  charcoal  burns 
and  greedily  takes  all  the  oxygen  it  can  get,  and  produces 
carbon  dioxide.  But  as  this  passes  along  the  red  hot  tube  it 
gives  up  half  of  its  oxygen  to  a  further  portion  of  carbon,  so 
that  they  go  equal  shares,  as  it  were,  with  the  oxygen,  and  the 
result  is  carbon  monoxide. 

C02  +  C  =  2CO. 

We  can  prove  that  this  is  the  true  explanation  in  the  follow- 
ing way. 

Experiment  209. — Take  a  piece  of  ordinary  iron  gas-pipe  and 
fill  it  with  fragments  of  charcoal  and  make  it  red  hot  in  a  furnace 
(Fig.  99).  (If  more  convenient,  the  iron  pipe  may  be  bent  in  the 
middle  into  a  sort  of  elbow  shape,  and  the  bent  part  pushed  into  an 
ordinary  fire,  between  the  bars.) 

Now  we  must  not  pass  oxygen  direct  into  this  pipe,  or  the  iron 
itself  would  burn  up  (refer  back  to  oxygen)  ;  so  we  must  connect  to 
one  end  of  it  a  short  piece  of  glass  combustion  tube  containing 


232  Oxides  of  Carbon. 

a  few  pieces  of  charcoal.  Now  heat  the  charcoal  in  the  glass  tube 
and  pass  a  gentle  stream  of  oxygen  through.  The  carbon  burns, 
and  the  carbon  dioxide  passes  along  over  the  red  hot  charcoal  in  the 
iron  pipe.  Set  fire  to  the  gas  which  escapes  at  the  other  end,  and 
observe  its  flame.  This  gas  is  carbon  monoxide. 

Carbon  monoxide  is  always  being  produced  by  exactly  this 
method  in  our  ordinary  fireplaces.  The  oxygen  of  the  air 
which  enters  the  front  and  bottom  of  the  grate  is  taken  up  by 
the  first  portion  of  burning  coal,  and  forms  with  the  carbon 
carbon  dioxide.  This,  in  passing  through  the  fire,  meets  with 
red-hot  carbon,  and  gives  up  half  its  oxygen,  forming  carbon 


FIG.  99. 

monoxide.  This  operation  always  goes  on  in  the  fire,  but  to 
a  greater  extent  if  the  fire  is  hot  and  clear,  and  especially  in  a 
coke  fire. 

A  good  deal  of  the  carbon  monoxide  thus  formed  burns  on 
the  top  of  the  fire,  and  the  familiar  bluish  flame  seen  flickering 
about  on  the  top  of  a  clear  bright  fire  is  this  gas  burning.  But 
some  of  it  escapes  unburnt  up  the  chimney,  because  there  is 
often  so  much  carbon  dioxide  passing  off  from  the  fire  as  to 
prevent  the  other  gas  from  burning. 

If  instead  of  passing  carbon  dioxide  through  the  tube  of 
red-hot  charcoal,  we  were  to  send  steam  through  it,  we  should 
then  get  a  mixture  of  carbon  monoxide  and  hydrogen, 

C  +  H2O  =  CO  -f  H* 
This  mixture  is  sometimes  called  water-gas,  and  is  made 


Carbon  Monoxide.  233 

on  a  large  scale  by  passing  steam  over  strongly  heated  coal 
or  coke. 

Preparation  of  Carbon  Monoxide. 

Experiment  210. — Put  a  few  crystals  of  oxalic  acid  into  a  test- 
tube,  and  pour  upon  them  a  little  strong  sulphuric  acid.  Fit  a  cork 
and  delivery  tube  to  the  test-tube  and  gently  heat  the  mixture. 
Notice  that  a  brisk  effervescence  soon  begins  to  take  place.  Collect 
two  jars  of  the  gas,  over  water.  Test  the  gas  in  one  jar  with  a 
lighted  taper,  and  note  that  it  burns  with  a  blue  flame,  but  does  not 
seem  to  burn  very  well. 

Into  the  other  jar  pour  some  lime  water  ;  notice  that  milkiness 
is   at   once  produced.      This   shows   that   carbon 
dioxide  is  also  present,  because  no  other  gas  pro- 
duces this  turbidity  with  lime  water. 

To  find  out  how  much  carbon  dioxide  is 
present,  we  can  make  use  of  the  fact  that  this 
gas  is  quickly  absorbed  by  caustic  soda. 

Experiment  211. — Collect  some  more  of  the 
gas  given  off  by  the  oxalic  acid  in  a  burette  (fill 
the  burette  with  water,  and  invert  it  in  a  basin  or 
trough,  just   as   an   ordinary  gas-collecting  jar). 
When  it  is  full  of  gas,  attach  a  small  funnel  to  the 
top  with  a  piece  of  rubber  tube,  as  shown  in 
Fig.  ioo.     Pour  some  caustic  soda  solution  into 
the  funnel,  and  then  gradually  turn  the  tap  so  as 
to  allow  the  liquid  slowly  to  enter  the  tube  and 
trickle  down  the  side.      Close  the  tap  before  the 
funnel  is   empty,  or  else   air  will    be   drawn  in.          FIG  i< 
Notice  that  as  the  caustic  soda  enters,  the  gas  is 
quickly  absorbed  by  it,  and  the  water  rises  in  the  tube.     When 
it   stops  rising,  note    how    much    gas    remains  ;    almost   exactly 
one  half. 

Close  the  tube  with  the  thumb,  remove  it  from  the  trough  and 
invert  it.  Then  bring  a  lighted  taper  to  the  gas.  Notice  that  it 
burns  with  the  blue  flame  of  carbon  monoxide. 

This  shows  that,  when  oxalic  acid  is  acted  on  by  sulphuric 
acid,  carbon  dioxide  and  monoxide  are  given  off  in  equal 
volumes.  The  equation  expressing  the  change  is  this — 

C2H2O4  =  C02  +  CO  +  H20. 

Oxalic  acid. 


234  Oxides  of  Carbon. 

The  water  is  taken  up  by  the  sulphuric  acid,  which  has  a 
powerful  affinity  for  water. 

When  we  want  carbon  monoxide  pure,  we  must  pass  the 
gas  obtained  from  oxalic  acid  through  bottles  containing 
caustic  soda,  so  as  to  get  the  carbon  dioxide  removed. 

Experiment  212. — Fit  up  the  apparatus  shown  in  Fig.  101. 
Place  in  each  of  the  two  bottles  some  solution  of  caustic  soda, 

and  heat  the  mixture  of 
oxalic  acid  and  sulphuric 
acid  in  the  small  flask.  As 
the  gases  bubble  through 
the  caustic  soda,  the  carbon 
dioxide  is  absorbed,  and 
the  monoxide  passes  on. 
Collect  three  jars. 

Dip  a  lighted  taper  into 

FIG  IQI  one.     Notice  that  the  gas 

burns  with  a  much  stronger 

flame  than  before.  '  Also  note  that  the  taper  itself  is  extinguished 
if  thrust  into  the  gas. 

Add  lime  water  to  the  second.  Is  there  any  turbidity  produced  ? 
If  so,  it  proves  that  even  bubbling  through  two  bottles  has  not 
quite  removed  all  the  carbon  dioxide.  Take  the  third  jar,  and 
quickly  pour  into  it  a  little  caustic  soda  and  cover  it  again  imme- 
diately. Shake  the  liquid  up  with  the  gas,  and  then  replace  the 
jar  in  the  trough  for  a  few  minutes,  in  order  to  let  the  caustic  soda 
go  out  into  the  water.  Now  take  the  jar  out  again,  and  pour  some 
lime  water  into  it  and  shake  up.  This  time  there  should  be  no 
turbidity  at  all.  Now  light  the  gas,  and  as  the  flame  passes  down 
into  the  jar,  cover  it  with  the  glass  plate.  Again  shake  up  the 
lime  water  that  is  in  the  jar,  and  notice  that  now  it  is  instantly 
made  milky. 

This  shows  that  when  carbon  monoxide  burns,  it  gives 
carbon  dioxide. 

CO  +  O  =  C02. 

Carbon  monoxide  takes  oxygen  away  from  many  metallic 
oxides  when  they  are  strongly  heated  in  this  gas.  It  is,  there- 
fore, like  carbon,  a  reducing  agent.  Thus,  if  carbon  monoxide 
is  passed  over  heated  oxide  of  iron,  the  oxide  is  deprived  of 


Carbon  Monoxide.  235 

its  oxygen  and  becomes  reduced  to  the  metallic  state,   and 
carbon  dioxide  is  formed. 

Fe2O3  +  SCO  =  3CO2  +  2Fe. 

This  process  goes  on  in  the  blast  furnace  where  iron  ores 
are  smelted. 

Carbon  monoxide  is  extremely  poisonous.  Many  people 
have  been  killed  by  the  gas  escaping  from  coke  or  charcoal 
fires,  their  deaths  being  due  to  the  poisonous  nature  of  the 
carbon  monoxide  which  such  fires  give  off. 

EPITOME. 

Carbon  monoxide  is  produced  when  carbon  burns  in  an  insuffi- 
cient supply  of  oxygen,  or  when  carbon  dioxide  is  passed  over  red 
hot  charcoal. 

It  is  prepared  from  oxalic  acid  by  the  action  of  sulphuric  acid, 
the  gas  so  obtained  being  passed  through  caustic  soda  to  remove 
the  carbon  dioxide.  Carbon  monoxide  is  a  colourless  poisonous  gas, 
which  burns  with  a  beautiful  blue  flame.  When  it  burns  it  produces 
carbon  dioxide. 

Carbon  monoxide  does  not  dissolve  in  water ;  does  not  make 
lime  water  turbid  ;  is  not  absorbed  by  caustic  potash.  It  is  not  an 
acid  forming  oxide  like  carbon  dioxide,  and,  therefore,  forms  no  salts. 

It  is  readily  distinguished  from  all  other  gases  by  burning  with 
a  blue  flame,  and  forming  carbon  dioxide. 


CHAPTER  XXVI. 

SULPHUR. 

THIS  element  is  found  chiefly  in  volcanic  regions,  such  as 
Sicily  and  Iceland,  where  it  exists  in  the  free  state ;  that  is, 
not  in  chemical  combination  with  other  elements.  The 
sulphur  as  thus  found  is  called  native  sulphur. 

Besides  occurring  in  this  uncombined  state,  sulphur  is  also 
a  constituent  of  a  large  number  of  important  ores,  in  which  it 
is  combined  with  various  metals.  Some  of  the  commonest  of 
these  sulphides  are  iron  pyrites  (sulphur  combined  with  iron, 
FeS2) ;  copper  pyrites  (sulphur  with  copper  and  iron,  CuFeS2)  ; 
zinc  blende  (sulphur  and  zinc,  ZnS)  ;  galena  (sulphur  and  lead, 
PbS). 

In  combination  with  metals  and  with  oxygen  together,  it 
is  found  in  heavy  spar  (barium  sulphate,  BaSO4) ;  and  the  very 
common  mineral  gypsum  (calcium  sulphate,  CaSO4  +  H2O). 

Modes  of  obtaining  Sulphur. — (i)  "Native  sulphur" 
is  always  mixed  up  with  more  or  less  earthy  and  mineral 
matters.  In  order  to  separate  the  sulphur,  the  crude  material 
is  piled  up  into  heaps  on  a  slanting  hearth,  and  the  heaps  set 
on  fire.  Some  of  the  sulphur  burns  away,  but  the  heat  it  gives 
out  melts  the  remainder,  which  runs  away  from  the  impurities 
down  the  sloping  floor.  Of  course  the  supply  of  air  to  the 
heap  requires  to  be  regulated,  for  if  it  had  free  access,  the 
whole  of  the  sulphur  would  be  burnt  away.  (2)  Sulphur  can 
also  be  got  from  iron  pyrites,  by  heating  the  ore  strongly, 
without  letting  air  get  to  it. 

Experiment  213.     Heat  a  little  powdered  iron  pyrites  in  a  test- 


Purification  of  Sulphur. 


237 


tube.  Notice  the  sublimate  which  collects  on  the  cooler  part  of  the 
tube,  as  it  forms  little  drops  of  melted  sulphur.  After  a  time,  just 
touch  the  hot  end  of  the  test-tube  with  a  drop  of  water  so  as  to 
crack  it  off.  Now  heat  the  sulphur  on  the  side  of  the  tube,  holding 
the  tube  in  an  inclined  position.  Notice  the  sublimed  sulphur 
takes  fire  and  burns,  and  the  gas  which  escapes  at  the-  top  of  the 
tube  has  the  characteristic  choking  smell  produced  by  burning 
sulphur. 

The  pyrites  does  not  part  with  all  its  sulphur  when  heated 
in  this  way.     The  change  is  expressed  by  the  equation— 
3FeS2  =  Fe3S4  +  28. 

If  the  iron  pyrites  is  roasted  in  a  free  current  of  air,  it  loses 
all  its  sulphur,  and  both  the  iron  and  the  sulphur  are  con- 
verted into  oxides,  thus  — 

2FeS3  -f-  nO  =  Fe2O3 


A  large  quantity  of  sulphur  is  now  recovered  from  waste 
products  which  contain  this  element,  and  which  used  to  be 
thrown  away.  Such  a  refuse  sub- 
stance as  that  known  as  Alkali- 
waste  (obtained  in  the  process  of 
manufacturing  sodium  carbonate) 
is  now  utilized  in  this  way. 

Purification  of  Sulphur. 

Experiment  214.  —  Bend  a  piece  of 
wide  glass  tube,  one  end  of  which  is 
closed  up,  and  attach  it  to  a  retort  in 
the  manner  shown  in  Fig.  102.  Place 
a  few  pieces  of  sulphur  in  the  little 
bent  tube,  and  apply  heat  to  it.  The 
sulphur  melts  ;  then  gets  very  dark 
in  colour,  and  presently  boils,  and  the 
vapour  passes  into  the  body  of  the 
retort.  The  sulphur  is  here  being 
distilled,  and  the  impurities  it  may 
contain  remain  behind.  Notice  that, 
as  the  vapour  enters  the  large  retort,  some  of  it  condenses  upon  the 
glass  as  a  yellowish  powder,  while  some  collects  on  the  lower  part 
of  the  vessel  in  the  liquid  state  and  quickly  solidifies. 


FIG.  102. 


238  Sulphur. 

On  a  large  scale  this  process  is  carried  out  by  boiling  the 
sulphur  in  earthenware  retorts,  and  sending  the  vapour  into 
large  brickwork  chambers.  At  first  it  condenses  as  a  fine 
light  yellow  powder  on  the  walls.  In  this  condition  it  is 
called  flowers  of  sulphur.  After  a  time  the  walls  of  the 
chamber  get  warm,  and  the  sulphur  melts  and  collects  on 
the  floor,  and  is  then  run  out  into  wooden  moulds  so  as  to 
cast  it  in  the  form  of  sticks.  This  is  called  roll  sulphur,  or 
brimstone. 

Properties  of  Sulphur. 

Experiment  215. — Take  a  piece  of  common  roll  sulphur,  and 
strike  it  gently  on  the  table,  or  with  a  pestle.  Notice  how  brittle 
it  is.  Examine  the  freshly  broken  surfaces  and  see  that  it  is  highly 
crystalline.  Powder  a  little  of  it  in  a  mortar,  and  try  if  it  dissolves 
in  water.  After  shaking  it  up  with  water  in  a  test-tube  for  a  little 
time,  decant  some  of  the  water  into  a  small  dish  and  evaporate 
gently  to  dry  ness.  If  there  is  nothing  left  it  shows  that  sulphur  is 
not  dissolved  by  water. 

When  sulphur  is  heated  it  behaves  in  a  rather  striking 
manner. 

Experiment  216. — Carefully  heat  a  little  of  the  powdered  sulphur 
in  a  test-tube ;  it  easily  melts,  and,  if  not  over-heated,  gives  a  pale 
amber-coloured  liquid  which  runs  about  the  tube  like  oil.  Now 
heat  more  strongly,  and  notice  that  the  sulphur  rapidly  deepens  in 
colour,  becoming  like  dark  treacle,  and  gets  so  thick  and  sticky 
that  if  the  tube  is  turned  upside  down  it  does  not  run  out  at  all. 
Heat  it  still  more,  and  note  that  it  becomes  quite  liquid  again, 
although  remaining  dark  coloured,  and  presently  boils.  It  is 
difficult  to  see  the  colour  of  the  vapour,  because  of  the  almost 
black  appearance  of  the  liquid  ;  but  by  looking  through  the  tube 
at  any  part  where  the  glass  is  clear,  it  will  be  seen  that  the  vapour 
has  a  pale  yellow  colour.  Let  the  test-tube  cool,  and  the  sulphur 
goes  through  the  same  changes  in  the  opposite  order.  Notice  that 
as  it  solidifies  it  forms  crystals  on  the  sides  of  the  tube. 

Sulphur  melts  at  114*5°;  at  a  temperature  about  230°  it 
passes  into  the  thick  condition,  and  at  448°  it  boils. 

Allo tropic  Modifications. — Sulphur,  like  carbon,  exists 
in  three  allotropic  forms;  and,  like  those  of  carbon,  two  are 


Allotropic  Forms  of  Sulphur. 


239 


crystalline,  and  one  non-crystalline  or  amorphous.     But  here 
the  similarity  ends. 

Sulphur  is  quite  easily  made  to  assume  either  of  its  three 
forms  (not  so,  carbon),  but  it  will  only  remain  in  one  of  them, 
for  both  the  other  varieties  gradually  pass  back  into  the  first. 
(Carbon  is  stable  in  each  of  its  allotropic  forms.) 

Experiment  217. — Put  some  fragments  of  roll  sulphur  into  a 
test-tube,  and  add  to  them  a  small  quantity  of  a  liquid  called 
carbon  disulphide  (CS2)  just  to  cover  them.  Notice  that  the 
sulphur  quickly  dissolves.  Pour  the  solution  into  a  small  dish, 
cover  it  .with  a  piece  of  cardboard,  and  leave  it  for  some  time  to* 
slowly  evaporate.  Examine  the  residue  with  a  pocket  lens,  and 
carefully  note  the  shapes  of  the  crystals. 

If  this  experiment  is  made  on  a  larger  scale,  and  with 
certain  precautions,  more  perfectly  shaped  crystals  will  be 
obtained,  like  the  one  shown 
in  Fig.  103. 

The  form  of  the  crystals 
of  sulphur  obtained  in  this 
way  is  what  is  known  as 
Rhombic  Octahedral.  They 
have  a  beautiful  amber-like 


and     are 


very 


appearance, 
brittle. 

The  sulphur  which  is 
found  "native"  is  in  this 
form. 


FIG.  103. 


Experiment  2 1 8.— Carefully  melt  some  roll  sulphur  in  a  small 
beaker;  the  beaker  being  about  three-quarters  full.  Allow  it  to 
cool,  carefully  watching  it,  and  as  soon  as  a  thin 
crust  has  formed  on  the  top,  pour  out  what 
remains  of  the  liquid  into  a  dish  or  plate.  On 
cutting  away  the  crust,  the  interior  of  the  beaker 
will  be  found  to  be  lined  with  long  needle-shaped 
crystals,  like  those  shown  in  Fig.  104.  Examine 
these  crystals,  and  note  how  entirely  different  they 
are  in  shape  to  the  others.  They  have  also  a 
transparent  appearance. 


FIG.  104. 


240  Sulphur. 

This  is  the  second  allotropic  form  of  sulphur,  and  is  known 
by  the  name  Prismatic  Sulphur,  because  the  crystals  are  in 
the  shape  of  long  thin  prisms.  If  these  crystals  are  kept  for 
a  day  or  two  they  lose  their  transparent  appearance,  and 
become  exactly  the  colour  of  ordinary  brimstone.  In  fact, 
they  change  back  again  from  the  prismatic  modification  to 
the  rhombic  form;  and  although  the  crystals  retain  the  out- 
ward form  of  the  prism,  they  crumble  on  the  slightest  touch 
to  a  number  of  minute  crystals,  having  the  shape  of  rhombic 
octahedrons. 

Experiment  219. — Heat  a  quantity  of  sulphur  (either  "flowers" 
or  powdered  lump)  in  a  common  oil  flask  until  it  gets  into  a  boiling 
condition.  Then  pour  the  hot  sulphur  in  a 
thin  stream  into  cold  water  in  a  beaker. 
(A  funnel  may  be  stood  in  the  water,  and  the 
stream  of  sulphur  poured  round  and  round 
it,  Fig.  105.)  Now  lift  out  the  funnel  with 
the  congealed  sulphur,  and  notice  how  en- 
tirely different  it  is  from  either  of  the  other 
forms.  It  is  no  longer  brittle,  but  seems 
almost  like  indiarubber,  and  can  be  stretched 
"FIG!  105.  and  pulled  into  threads. 

This  is  the  third  allotrope  of  sulphur,  and  is  called  Plastic 
Sulphur.  It  has  no  crystalline  character  at  all. 

Like  the  prismatic  form,  this  plastic  sulphur  when  left  to 
itself  for  a  few  days  changes  back  again  into  the  rhombic 
variety.  It  gradually  loses  its  curious  elasticity,  and  becomes 
the  ordinary  brittle,  yellow,  crystalline  sulphur.  If  it  is 
stretched  about,  or  slightly  warmed,  it  changes  from  its  plastic 
state  to  the  ordinary  condition  much  more  quickly. 

Milk  of  Sulphur  is  an  old-fashioned  name  for  sulphur 
obtained  in  the  following  way. 

Experiment  220. — Throw  a  small  handful  of  flowers  of  sulphur 
into  a  saucepan  half  full  of  hot  water,  and  add  about  twice  as 
much  lime.  Boil  the  mixture  for  five  or  ten  minutes,  and  then 
allow  it  to  settle.  Take  some  of  the  clear  yellowish  liquid  and 
add  to  it  a  little  strong  hydrochloric  acid.  A  white  precipitate  of 
sulphur  in  the  condition  of  very  fine  particles  is  produced,  which 


Sulphides  241 

makes  the  liquid  look  almost  like  milk  ;  hence  the  name  "  Milk 
of  Sulphur." 

This  is  a  much  finer  powder  than  flowers  of  sulphur,  and 
is  on  this  account  used  in  medicine. 

When  sulphur,  in  any  of  its  allotropic  forms,  is  heated  in 
air,  it  burns  with  a  blue  flame  (see  Exps.  64,  138),  and  gives 
an  oxide  of  sulphur  called  sulphur  dioxide,  SO2. 

Combination  of  Sulphur  with  Metals.— Sulphur 
combines  with  many  metals,  and  forms  compounds  called 
sulphides  ;  just  as  oxygen  unites  with  metals  and  gives  oxides. 

Experiment  22 1 . — Heat  a  small  quantity  of  flowers  of  sulphur 
in  a  test-tube  until  it  boils,  and  drop  in  a  fragment  of  sodium  about 
the  size  of  a  large  pin's  head.  The  sodium  instantly  takes  fire  and 
burns  brilliantly,  producing  sodium  sulphide,  Na2S. 

Experiment  222. — Drop  into  a  similar  small  quantity  of  boiling 
sulphur  some  reduced  iron  (that  is,  iron  obtained  by  heating  iron 
oxide  in  a  stream  of  hydrogen  or  coal  gas).  The  black  powder  at 
once  takes  fire  as  it  comes  into  the  sulphur  vapour,  and  gives  ferrous 
sulphide,  FeS.  If  iron  filings  are  used  instead  of  the  "reduced" 
iron,  they  combine  with  the  sulphur,  but  being  so  much  coarser 
they  do  not  take  fire  ;  but  if  a  fine  iron  wire  is  made  red  hot  in  a 
Bunsen  flame,  and  then  plunged  into  the  sulphur  vapour,  the  wire 
burns,  and  melted  drops  of  iron  sulphide  fall  to  the  bottom  of 
the  tube. 

Thin  copper  wire  also  readily  combines  with  sulphur  vapour, 
and  takes  fire  without  being  heated  in  a  lamp,  producing  copper 
sulphide,  CuS. 

EPITOME. 

Sulphur  occurs  in  volcanic  regions,  such  as  Sicily  and  Iceland, 
in  the  free  or  elementary  state.  In  combination  with  many  metals, 
in  some  of  the  commonest  ores  of  these  metals,  as  sulphides ; 
also  in  combination  with  metals  and  oxygen,  as  sulphates. 

"  Native  sulphur  "  is  burnt  in  heaps,  with  a  limited  air  supply, 
so  as  to  melt  the  sulphur  and  separate  it  from  rocky  matter  with 
which  it  is  mixed. 

Sulphur  is  got  from  iron  pyrites  by  simply  heating  it,  when  it 
gives  off  one-third  of  its  sulphur. 

Sulphur  is  purified  by  distillation ;  and  is  obtained  either  as 
flowers,  or  is  melted  and  cast  into  sticks,  known  as  roll  sulphur. 

R 


242  Sulphur. 

The  three  allotropic  forms  of  sulphur  are — 

(1)  Rhombic    (octahedral).      Permanent,    soluble    in    carbon 
disulphide.     Specific    gravity,    2^05 .     Lemon   yellow   colour,   very 
brittle. 

(2)  Prismatic.     Not    permanent,   slowly    changes   to    No.    I. 
Translucent  amber  yellow  crystals.     Specific  gravity,  1*98. 

(3)  Plastic.     Not  permanent,  slowly  changes  to  No.  i.     Soft, 
non-crystalline,    indiarubber-like,    translucent    yellow.      Not    dis- 
solved by  carbon  disulphide.     Specific  gravity,  1-95. 

Sulphur  (any  variety)  burns  in  the  air  with  a  blue  flame, 
producing  sulphur  dioxide. 

Sulphur  is  not  soluble  in  water,  but  is  oxidized  by  nitric  acid 
into  sulphuric  acid  (see  p.  191).  It  unites  directly  with  metals, 
forming  sulphides. 

Sulphur  belongs  to  a  family  of  elements,  the  members  of  which 
are  oxygen,  sulphur,  selenium,  and  tellurium. 

The  last  two  are  rather  rare  elements. 

The  chemical  relationship  between  sulphur  and  oxygen  will 
be  seen  by  comparing  some  of  the  compounds  which  each 
forms  with  other  elements. 

Water,  H2O  Sulphuretted  hydrogen,  H2S. 

Potassium  hydroxide,  KHO  Potassium  hydrosulphide,  KHS. 

Calcium  hydroxide,  Ca(HO)2  Calcium  hydrosulphide,  Ca(HS)2. 

Carbon  dioxide,  CO2  Carbon  disulphide,  CS2. 

Potassium,  oxide,  K2O  Potassium  sulphide,  K2S. 

Copper  oxide,  CuO  Copper  sulphide,  CuS. 


CHAPTER   XXVII. 

SULPHURETTED    HYDROGEN,    H2S. 

Occurrence. — This  compound  is  a  gas,  and  is  present  in 
volcanic  gases.  It  is  met  with,  dissolved  in  water,  in  certain 
sulphur  springs,  such  as  those  at  Harrowgate.  When  animal 
substances  containing  sulphur  become  putrid,  sulphuretted 
hydrogen  is  formed;  bad  eggs  owe  their  disagreeable  smell 
to  the  presence  of  this  gas.  Ordinary  coal-gas,  as  it  leaves 
the  retorts,  contains  a  considerable  amount  of  sulphuretted 
hydrogen.  This  is  removed  in  the  "  purifiers  "  before  the  gas 
is  sent  out  from  the  gasworks,  and  the  sulphur  it  contains  is 
extracted  and  sold. 

Modes  of  Formation. — Sulphur  does  not  very  easily 
combine  with  hydrogen ;  but  if  hydrogen  is  passed  over  boiling 
sulphur,  a  portion  of  the  sulphur  unites  with  hydrogen  and 
forms  sulphuretted  hydrogen. 

Experiment  225. — Heat  a  few  fragments  of  sulphur  in  a 
horizontal  bulb  tube,  until  the  sulphur  boils,  and  then  allow 
hydrogen  to  pass  through  the  tube.  Smell  the  gas  escaping  at  the 
end.  Also  test  it  by  holding  a  piece  of  paper  moistened  with  a 
solution  of  lead  acetate  in  the  gas.  The  paper  becomes  black, 
showing  the  presence  of  sulphuretted  hydrogen. 

In  the  laboratory,  the  gas  is  always  prepared  by  another 
method,  namely,  by  acting  on  ferrous  sulphide  with  either 
sulphuric  or  hydrochloric  acid.  Ferrous  sulphide  is  made 
by  heating  iron  and  sulphur  together.  Its  composition  is 
expressed  by  the  formula  FeS.  This  substance  must  not  be 
confounded  with  iron  pyrites,  FeS2. 


244  Sulphuretted  Hydrogen. 

Experiment  224. — Place  a  quantity  of  ferrous  sulphide  in  a 
two-necked  bottle  (see  Fig.  39),  cover  it  with  water,  and  pour 
a  small  quantity  of  strong  sulphuric  acid  through  the  thistle  funnel. 
Notice  that  effervescence  immediately  begins.  After  a  few  minutes, 
during  which  the  air  in  the  bottle  is  being  gradually  swept  out, 
collect  two  jars  of  the  gas,  using  water  in  the  trough  as  warm  as 
the  hands  can  comfortably  bear. 

The  reaction  in  this  case  is — 

FeS       +      H2SO4      =       FeSO4      +      H2S. 

Ferrous  sulphate. 

If  hydrochloric  acid  is  used  instead  of  sulphuric  acid  the 
equation  is — 

FeS      +      2HC1      =       FeCl2      +      H2S. 

Ferrous  chloride. 

Then  remove  the  delivery  tube  from  the  apparatus  and  light 
the  gas  as  it  escapes  from  the  exit  tube.  Notice  that  the  flame  is 
bluer  than  that  of  hydrogen,  but  not  so  blue  as  that  of  burning 
sulphur.  Gently  smell  the  products  of  the  burning  gas  ;  note  that 
there  is  the  same  choking  smell  of  sulphur  dioxide,  as  when  sulphur 
burns.  Hold  a  cold  tumbler  over  the  flame,  and  observe  the 
moisture  collecting. 

When  sulphuretted  hydrogen  burns  from  a  jet  with  a  free 
supply  of  air,  it  gives  sulphur  dioxide  and  water  - 

H2S  +  30  =  S02  +  H20. 

Experiment  22$. — Depress  a  piece  of  glass  down  on  to  the 
flame,  and  note  that  there  is  a  deposit  of  sulphur. 

Experiment  226. — Test  the  gas  in  one  of  the  jars  with  a  lighted 
taper,  note  that  the  taper  is  extinguished  when  thrust  into  the  gas. 
Observe  also  that,  as  the  gas  burns,  there  is  a  deposit  of  sulphur 
formed  on  the  sides  of  the  jar.  This  is  from  the  same  cause 'as  in 
the  last  experiment,  namely,  because  the  supply  of  air  is  not 
sufficient  to  completely  burn  the  gas.  We  may  express  it  by  this 
equation — 

H2S  +  O  =  H2O  +  S. 

Experiment  227. — Transfer  the  second  jar  of  gas  to  a  trough 
of  cold  water,  and  leave  it  standing  mouth  downwards.  Notice 
that  the  water  rises  a  little  in  the  jar.  Shake  it  up  as  much  as 
possible  without  lifting  the  mouth  out  of  the  water,  and  in  a  short 


Sulphuretted  Hydrogen.  245 

time  the  water  will  have  absorbed  nearly  all  the  gas.     Smell  the 
water ;  note  that  it  smells  like  the  gas. 

Sulphuretted  hydrogen  is  considerably  soluble  in  cold 
water.  At  the  common  temperature,  water  dissolves  about 
three  times  its  own  bulk  of  this  gas.  The  solution  is  called 
sulphuretted  hydrogen  water.  Warm  water  dissolves  much 
less  of  the  gas,  therefore  we  usually  collect  it  over  hot  water, 
as  described  in  Exp.  224. 

Sulphuretted  hydrogen  made  from  ferrous  sulphide  is  always 
mixed  with  free  hydrogen,  because  the  ferrous  sulphide  always 
contains  some  iron  which  is  not  combined  with  sulphur  ;  and 
when  the  acid  comes  in  contact  with  this  free  iron,  hydrogen 
•is  evolved  (see  Hydrogen,  p.  46).  Therefore,  if  sulphuretted 
hydrogen  is  required  quite  pure,  we  generally  use  antimony 
sulphide,  Sb2S3.  For  the  purposes  for  which  the  gas  is  generally 
required,  however,  the  presence  of  a  little  hydrogen  does  not 
matter,  so  that  in  practice  it  is  nearly  always  made  from  the  iron 
compound.  For  a  number  of  experiments  we  want  just  a  few 
bubbles  of  sulphuretted  hydrogen,  and  therefore  it  is  convenient 
to  have  an  apparatus  so  arranged  for  making  the  gas  that  we 
can  stop  the  action  when  we  please,  and  start  it  again. 
This  cannot  be  done  with  the  form  of  apparatus  used  for 
Exp.  224,  without  emptying  the  bottle  each  time.  A  very 
simple  form  of  constant  apparatus  can  be  made  in  the  following 
way. 

Experiment  228. — Obtain  two  large  test-tubes  ("  boiling  tubes  ") 
and  draw  them  out  at  one  end,  as  shown  in  Fig.  106.  Secure  one 
of  them  with  wire  or  thread  to  a  retort  stand,  and  join  their  drawn- 
out  ends  with  a  piece  of  indiarubber  pipe  in  the  manner  shown. 
Half  fill  the  fixed  one  with  small  broken  pieces  of  ferrous 
sulphide,  and  close  the  tube  with  a  cork  and  exit  tube,  the  latter 
carrying  a  short  piece  of  rubber  tube,  /,  with  a  screw  clamp,  s, 
upon  it. 

Suspend  the  other  boiling  tube  to  a  ring  by  means  of  a  wire 
hook.  Close  the  clamp,  and  pour  dilute  sulphuric  acid  into  the 
open  tube  until  it  is  about  three-quarters  full.  Now  gently  open  the 
clamp,  when  the  acid  will  gradually  enter  the  other  tube,  and, 
coming  in  contact  with  the  ferrous  sulphide,  will  cause  the  evolution 


246 


Sulphuretted  Hydrogen. 


of  sulphuretted  hydrogen.  So  long  as  the  clamp  is  open,  the  gas 
will  escape  from  the  tube,  but  as  soon  as  it  is  closed  again,  the  gas, 
which  is  still  being  produced,  not  being 
able  to  get  out,  begins  to  drive  the  acid 
back  into  the  open  tube,  and  then,  of 
course,  the  action  stops.  In  this  way, 
by  opening  and  closing  the  clamp,  we 
can  produce  a  little  gas  and  then  stop 
the  supply  at  will. 

By  means  of  this  little  apparatus 
for  getting  sulphuretted  hydrogen,  do 
the  following  experiments — 

Action  of  Sulphuretted  Hy- 
drogen on  Metals. 

Experiment  229.— Open  the  screw 
clump,  s,  and  let  the  gas  blow  against  a 
clean  silver  coin  for  a  moment.  Notice 
that  the  silver  is  at  once  turned  black. 
[A  silver  spoon  is  stained  in  the  same 
way  by  a  bad  egg.]  The  black  substance 
is  silver  sulphide,  Ag2S. 

Experiment  230. — Place  a  small  frag- 
FlG  Io6  ment  of  potassium  in  a  horizontal  bulb 

tube,  and  pass  sulphuretted  hydrogen 

through  the  tube.  Heat  the  potassium,  when  it  will  burn  brightly 
in  the  gas.  Sulphuretted  hydrogen,  therefore,  will  support  the 
combustion  of  potassium.  The  reaction  is — 

H2S     +     K     =     KHS     +     H. 

Potassium  hydrosulphide. 

Compare  this  with  the  action  of  potassium  upon  water — 
H2O  +  K=  KHO  +  H. 

Action    of   Sulphuretted    Hydrogen    on    Metallic 
Compounds. 

Experiment  231. — Let  sulphuretted  hydrogen  blow  against  a  piece 
of  ordinary  canvas,  used  by  artists.  Notice  that  it  is  quickly 
blackened.  The  canvas  is  painted  with  white  lead  (a  lead  car- 
bonate), this  is  at  once  converted  into  lead  sulphide,  which  is  black. 


Metallic  Sulphides.  247 

[There  is  always  a  little  sulphuretted  hydrogen  in  the  air  of  towns, 
and  this  makes  oil  paintings  gradually  turn  black.] 

Experiment  232. — Put  a  little  oxide  of  iron  in  a  horizontal  glass 
tube,  and  gently  warm  it.  Then  pass  sulphuretted  hydrogen 
through  the  tube.  Notice  that  the  reddish  oxide  is  turned  black, 
and  glows  with  the  heat  produced  by  its  combination  with  the 
sulphur — 

FeaOs  +  3H2S  =  2FeS  +  S  +  3HaO. 

Similarly  if  sulphuretted  hydrogen  is  passed  over  slaked 
lime  (calcium  hydroxide),  we  get  a  sulphur  compound  of 
calcium  produced. 

CaH2O2      +      2H2S       -       2H2O       +      CaH2S2. 

Calcium  hydroxide.  Calcium  hydrosulphide. 

These  two  substances,  namely,  slaked  lime  and  iron  oxide, 
are  the  materials  used  in  the  "  purifiers  "  of  the  gas-works,  to 
absorb  the  sulphuretted  hydrogen  from  the  coal-gas. 

Action  of  Sulphuretted  Hydrogen  on  Metallic 
Solutions. 

Experiment   233. —  Prepare  the   three  following    solutions  : — 

(1)  Dissolve  a  small  particle  of  lead  acetate  in  water  in  a  test-tube. 

(2)  Dissolve  a  similar  quantity  of  white  arsenic  (arsenious  oxide)  in 
a  few  drops  of  hydrochloric  acid,  and  add  water  to  half  fill  the  test- 
tube.     (3)  Dissolve  a  small  quantity  of  tartar  emetic  (an  antimony 
salt)  in  water. 

Now  pass  a  few  bubbles  of  sulphuretted  hydrogen  through  each 
of  these,  by  dipping  a  tube  from  the  apparatus  (Fig.  106)  into  the 
solutions  in  turn. 

Notice  what  happens  in  each  case.  We  get  a  black,  a  yellow, 
and  a  red  precipitate.  In  other  words,  the  three  elements,  lead, 
arsenic,  and  antimony,  form  sulphides  ;  lead  sulphide  being  black, 
arsenic  sulphide  yellow,  and  antimony  sulphide  red. 

Sulphuretted  hydrogen,  therefore,  affords  a  test  by  which 
we  can  easily  distinguish  between  compounds  of  these 
elements. 

Experiment  234. — In  three  separate  test-tubes  put  a  little  dilute 
solutions  of  (i)  copper  sulphate,  (2)  ferrous  sulphate,  and  (3)  potas- 
sium nitrate  ;  add  to  each,  one  or  two  drops  of  hydrochloric  acid. 


248  Sulphuretted  Hydrogen. 

and  pass  into  each  a  few  bubbles  of  sulphuretted  hydrogen.     Note 
that  there  is  a  precipitate  only  in  the  copper  solution. 

This  means  that  copper  sulphide  is  precipitated  in  an  acid 
solution ;  whereas  no  sulphide  of  either  iron  or  potassium  is 
produced  in  a  solution  containing  hydrochloric  acid. 

In  acid  solution,  CuSO4  +  H2S  =  CuS  +  H2SO4. 

Now  add  a  few  drops  of  ammonia  to  the  two  solutions  which 
were  not  precipitated,  and  note  that  in  the  case  of  the  iron  salt,  a 
black  precipitate  is  produced,  but  nothing  happens  in  the  solution 
of  the  potassium  salt. 

In  alkaline  solution,  FeSO4  +  H2S  =  FeS  +.H2SO4. 

This  black  precipitate  is  ferrous  sulphide  ;  and  this  experiment 
shows  that  in  an  alkaline  solution  iron  is  precipitated  in  the  form 
of  sulphide,  while  no  potassium  sulphide  is  produced  either  in  the 
acid  or  alkaline  solutions. 

The  following  experiment  shows  what  use  we  can  make  of 
these  facts — 

Experiment  235. — Pour  a  little  of  the  solutions  of  copper 
sulphate,  ferrous  sulphate,  and  potassium  nitrate  into  the  same 
test-tube.  Add  a  drop  or  two  of  hydrochloric  acid,  and  bubble 
sulphuretted  hydrogen  through  the  solution  for  a  few  minutes. 
We  know  by  the  former  experiment  that  only  the  copper  will  be 
precipitated  as  sulphide. 

Now  filter  the  liquid.  The  copper  sulphide  remains  on  the 
filter,  while  the  solution  which  passes  through  contains  the  iron  and 
potassium  salts.  Next  add  to  the  filtrate  some  ammonia.  This 
we  have  seen  causes  the  precipitation  of  ferrous  sulphide  (there 
being  sulphuretted  hydrogen  dissolved  in  the  solution).  Pass  this 
through  another  filter.  The  ferrous  sulphide  remains  on  the  filter 
and  the  potassium  salt  passes  through. 

In  this  way  we  have  separated  the  three  metals,  copper,  iron, 
and  potassium,  whose  salts  were  originally  mixed  in  the  solution. 

Sulphuretted  hydrogen  is,  therefore,  a  most  important 
agent  in  analysis;  for  we  find  (i)  a  certain  number  of  metals 
which  are  precipitated  as  sulphides  in  acid  solution,  (2)  others 
which  are  not  precipitated  as  sulphides  in  acid,  but  only  in 
alkaline  liquids,  and  (3)  others  which  are  not  precipitated  as 
sulphides  in  either  acid  or  alkaline  solutions. 


Test  for  Sulphuretted  Hydrogen.  249 

Besides  this,  as  we  have  seen,  some  sulphides  have  cha- 
racteristic colours,  by  which  they  are  easily  distinguished. 

Test  for  Sulphuretted  Hydrogen.— The  smell  of  this 
gas  is  sufficiently  characteristic  to  distinguish  it  from  all  others. 
An  additional  test  is  its  action  on  a  salt  of  lead,  such  as  lead 
acetate.  Paper  moistened  with  a  solution  of  lead  acetate  is 
turned  black  by  this  gas  owing  to  the  formation  of  lead 
sulphide. 

EPITOME. 

Sulphuretted  hydrogen  is  prepared  by  the  action  of  sulphuric 
or  hydrochloric  acid  on  ferrous  sulphide.  The  gas  has  a  dis- 
agreeable smell,  like  bad  eggs,  and  is  poisonous.  It  is  moderately 
soluble  in  cold  water,  giving  a  solution  having  the  smell  of  the  gas. 
It  must  be  collected  over  hot  water.  The  gas  burns  with  a  bluish 
flame,  giving  water  and  sulphur  dioxide  if  excess  of  air  is  present, 
but  depositing  some  of  its  sulphur  if  the  supply  of  air  is  limited. 

The  gas  combines  directly  with  many  metals,  as  lead,  copper, 
silver,  forming  sulphides.  The  brown  "  tarnish  "  which  comes  on 
silver  articles  exposed  to  the  air  of  towns,  is  due  to  the  formation 
of  silver  sulphide.  Sulphuretted  hydrogen  also  acts  on  metallic 
compounds,  both  in  the  solid  state  (as  in  the  case  of  "  white  lead," 
oxide  of  iron,  lime)  or  in  solution. 

On  account  of  its  behaviour  towards  metallic  salts  in  solution, 
it  is  used  in  analysis,  for  separating  and  detecting  the  various 
metals  whose  compounds  are  present.  Thus,  of  the  six  metals, 
lead,  copper,  iron,  zinc,  calcium,  potassium,  the  sulphides  of  lead 
and  copper 'are  precipitated  in  acid  solutions,  whilst  those  of  the 
others  are  not.  The  sulphides  of  iron  and  zinc  are  precipitated  in 
alkaline  solutions,  whilst  sulphides  of  calcium  and  potassium'  are 
soluble  in  both  acids  and  alkalies.  The  action  of  sulphuretted 
hydrogen  on  solutions  of  salts  of  these  metals  is  represented  by  the 
following  equations — 

Pb(NO3)2  +  H2S  =  PbS  +  2HNOS)    , 

CuS04       +  H2S  =  CuS  +  H2S04  }  In  acid  solutlOns' 

FeSO4       +  H2S  =  FeS  +  H2SO4  )    T 

ZnS04       +  H2S  =  ZnS  +  H2SO4      In  akalme  solutlons- 

CaCl2        +  H2S.     No  action. 

KC1  +  H2S.     No  action. 


CHAPTER  XXVIII. 

SULPHUR    DIOXIDE — SULPHUROUS    ACID — SULPHITES. 

THE  two  most  important  oxides  of  sulphur  are  sulphur  dioxide, 
SO2,  and  sulphur  trioxide,  SO3.  Both  of  these  compounds  are 
acid-forming  oxides.  The  first,  when  dissolved  in  water,  gives 
sulphurous  acid,  H2SO3;  whilst  the  second  gives  sulphuric 
acid,  H2SO4. 

Sulphur  Dioxide,  SO2. — This  compound  is  always 
produced  when  sulphur  is  burnt  in  the  air  or  in  oxygen,  so 
that  we  can  make  it  this  way — 

S  +  O2  =  SO* 

Experiment  236. — Place  a  piece  of  sulphur  in  a  short  horizontal 
piece  of  combustion  tube,  attach  a  delivery  tube  to  one  end,  and 
arrange  to  pass  oxygen  in  at  the  other.  Heat  the  sulphur  until  it 
begins  to  burn,  and  let  oxygen  slowly  stream  through.  Collect  the 
gaseous  product  by  downward  displacement,  covering  the  mouth 
of  the  cylinder  with  a  piece  of  paper  or  card.  [Note  that  the  gas 
is  not  clear.  This  is  because  there  is  always  produced  at  the  same 
time  a  small  quantity  of  sulphur  trioxide,  along  with  the  dioxide. ,] 
Test  the  gas  with  a  lighted  taper,  notice  that  the  gas  does  not  burn, 
and  that  it  puts  out  the  taper. 

Fan  a  little  of  the  gas  towards  the  face,  so  as  to  get  a  slight 
whiff  of  it.  Note  its  choking  smell,  familiarly  known  as  "  the  smell 
of  burning  sulphur." 

Enormous  quantities  of  sulphur  dioxide  are  made  in  this 
way  for  the  manufacture  of  sulphuric  acid.  Sometimes 
sulphur  itself,  and  sometimes  iron  pyrites,  is  used ;  and  it 
is  burnt,  not  in  glass  tubes  in  a  stream  of  pure  oxygen,  but  in 
special  furnaces  and  in  ordinary  air. 


Sulphur  Dioxide.  251 

Preparation.  —  When  we  want  sulphur  dioxide  for 
experiments  in  the  laboratory,  we  always  make  it  by  acting 
on  copper  with  sulphuric  acid. 

Experiment  237. — Place  a  quantity  of  scrap  copper  in  a  flask, 
fitted  with  a  cork  and  exit  tube,  and  pour  upon  it  enough  strong 
sulphuric  acid  to  well  cover  it.  Heat  the  acid  carefully  with  a  rose 
burner.  Notice  that,  as  the  acid  gets  warm,  it  becomes  muddy, 
owing  to  the  formation  of  a  black  powder,  which  collects  in  the 
flask  in  considerable  quantity.  Presently  effervescence  sets  in, 
and  sulphur  dioxide  is  rapidly  evolved.  Collect  four  jars  of  the  gas 
by  displacement. 

The  final  result  of  the  action  of  sulphuric  acid  on  copper 
is  expressed  by  the  equation — 

Cu  +  2H2S04  =  CuS04  +  2H2O  +  SO** 

Properties  of  Sulphur  Dioxide. — The  specimens  col- 
lected show  that  it  is  a  colourless  gas  :  and  its  smell  will  have 
been  perceived  during  the  preparation  of  it.  We  have  also 
seen,  by  Exp.  236,  that  the  gas  does  not  burn,  nor  support  the 
combustion  of  a  taper. 

Experiment  238. — Place  one  of  the  jars  of  gas  mouth  downwards 
in  water.  Notice  that  the  gas  is  absorbed  fairly  rapidly.  Add  a 
few  drops  of  litmus  to  the  water  and  note  that  the  solution  of  this 
gas  is  strongly  acid.  At  the  common  temperature,  water  dissolves 
about  fifty  times  its  own  volume  of  sulphur  dioxide. 

Experiment  239. — Pour  into  a  second  jar  of  the  gas  3  or  4  cc. 
of  litmus  solution,  and  shake  it  up.  As  before,  the  litmus  is  instantly 
reddened,  but  after  a  time  the  colour  gets  fainter  and  finally  almost 
entirely  disappears.  The  gas  has  bleaching  properties.  Now  add 
to  the  bleached  (or  nearly  bleached)  liquid  a  few  drops  of  strong 
sulphuric  acid.  Notice  that  the  red  colour  is  restored. 

Sulphur  dioxide  is  used  for  bleaching  straw,  flannel,  sponges, 
and  other  articles  which  would  be  injured  by  chlorine.  Its 

*  But  this  does  not  explain  everything  that  goes  on,  because  it  only 
tells  us  that  copper  sulphate,  water,  and  sulphur  dioxide  are  formed,  and 
takes  no  notice  of  the  black  substance  which  is  produced  in  the  flask. 
Copper  sulphate  is  not  black.  The  equation,  therefore,  only  gives  us  the 
final  products.  We  do  not  know  exactly  what  intermediate  products  are 
formed,  and  therefore  cannot  express  their  formation  by  an  equation. 


252 


Sulphur  Dioxide. 


bleaching  power  is  not  due  to  the  oxidation  of  the  colouring 
matter,  as  in  the  case  of  chlorine,  and  is  not  always  permanent. 
Sponges,  flannels,  and  straw  articles,  gradually  return  to  their 
unbleached  state. 

Flowers,  especially  those  of  a  violet  or  purplish  tint,  are 
quickly  bleached  if  placed  in  sulphur  dioxide. 

Although  sulphur  dioxide  does  not  support  the  combustion 
of  ordinary  combustibles,  some  things  will  burn  in  this  gas. 

Experiment  240. — Place  a   small   heap  of  lead  dioxide  on  a 
deflagrating  spoon,  gently  warm  it  and  then  lower  it  into  a  jar  of 
sulphur  dioxide.     The  two  dioxides  combine  with  so  much  energy 
that  the  lead  dioxide  becomes  red  hot,  and  forms  lead  sulphate. 
PbO2  +  SO2  =  PbSO4. 

Experiment  241. — Sprinkle  a  few  particles  of  sodium  dioxide 
(sodium  peroxide)  into  a  jar  of  sulphur  dioxide.  The  sodium 
dioxide  takes  fire  and  burns  brilliantly,  forming  sodium  sulphate, 

Na2O2  +  SO2  ==  Na2SO4. 

Sulphur  dioxide  is  used  as  a  disinfectant,  and  for  this 
purpose  is  generally  obtained  by  burning  sulphur. 


FIG.  107. 


When  sulphur  dioxide  is  cooled  below  —  8°  it  condenses 
to  a  liquid. 

Experiment  242. — Cut  a  test-tube  about  4  cm.  (nearly  2  inches) 
from  the  closed  end,  and  border  the  end.  Into  this  short  tube  fit 


Sulphites.  253 

a  cork,  carrying  two  long  thin  tubes,  as  shown  in  Fig.  107.  Place 
this  in  a  vessel  filled  with  a  mixture  of  powdered  ice  and  salt,  and 
connect  it  to  the  apparatus  for  making  sulphur  dioxide,  first  drying 
the  gas  by  bubbling  it  through  strong  sulphuric  acid  in  the  bottle  w. 
The  ice  and  salt  mixture  cools  the  tubes  below  —8°,  so  that  the  gas 
which  goes  through  will  be  condensed  to  the  liquid  state,  and  will 
collect  in  the  little  test-tube.  In  case  any  gas  passes  through  with- 
out condensing,  attach  a  delivery  tube  to  the  apparatus,  and  let  it  dip 
into  water.  This  will  absorb  any  escaping  gas.  When  enough 
liquid  is  collected,  lift  the  tube  out  of  the  freezing  mixture,  and  pour 
the  liquid  upon  a  little  water  in  a  small  dish.  The  sulphur  dioxide 
boiling  at  —8°  at  once  freezes  some  of  the  water.  Do  not  inhale  the 
gas,  as  it  is  very  irritating  to  the  lungs. 

Liquid  sulphur  dioxide  is  used  for  the  artificial  production  of 
ice  on  a  large  scale. 

Sulphurous  Acid,  H2SO3. — A  solution  of  this  acid 
is  produced  when  sulphur  dioxide  is  dissolved  in  water. 
The  solution  is  strongly  acid  towards  litmus,  and  smells  like 
the  gas.  When  it  is  warmed  the  gas  is  driven  off  again,  and 
the  whole  of  it  is  expelled  by  boiling  the  solution.  Sulphurous 
acid  does  not  keep.  It  slowly  absorbs  oxygen  from  the  air, 
and  changes  into  sulphuric  acid.  This  change  is  expressed  by 
the  equation — 

2S03  +  O  =  H2S04. 

On  account  of  the  readiness  with  which  it  takes  up  oxygen 
from  other  compounds,  sulphurous  acid  is  a  powerful  reducing 
substance. 

Experiment  243. — Dissolve  a  crystal  of  potassium  permanganate 
in  water.  This  substance  is  very  rich  in  oxygen  (KMnO4).  Pour 
this  deep  violet  solution  into  some  sulphurous  acid  (made  by  passing 
sulphur  dioxide  into  water).  Notice  that  the  colour  is  instantly 
discharged.  The  sulphurous  acid  takes  some  of  the  oxygen  away 
from  the  permanganate,  and  changes  into  sulphuric  acid. 

Sulphites  are  salts  of  sulphurous  acid,  obtained  by 
neutralizing  the  acid  with  a  base. 

Experiment  244. — Add  caustic  soda  solution  cautiously  to  some 
sulphurous  acid,  until  a  drop  of  the  liquid  on  a  glass  rod  just  turns 
reddened  litmus  paper  blue.  Then  add  one  or  two  drops  more  of 


254  Sulphur  Dioxide. 

the  acid  so  as  to  make  the  solution  just  acid.  Evaporate  the 
solution  to  dryness,  and  obtain  a  white  salt.  This  is  sodium 
sulphite. 

Add  a  few  drops  of  sulphuric  acid  to  it.  Notice  that  there  is 
effervescence.  Smell  the  gas,  and  note  that  it  is  sulphur  dioxide. 
Sulphuric  acid,  therefore,  decomposes  sulphites,  expelling  sulphur 
dioxide,  and  forming  sulphates.  The  reactions  here  are — 

(1)  H2S03    +  2NaHO  =  Na2SO3  +  2H2O. 

(2)  Na2SO3  +  H2SO4    =  Na2S04  +  H2O  +  SO2. 

Sulphurous  acid  contains  two  atoms  of  hydrogen  ;  it  is, 
therefore,  called  a  dibasic  acid.  It  can  form  two  classes  of 
salts  (just  as  carbonic  acid  does,  see  p.  227),  depending  on 
whether  both  or  only  one  of  these  hydrogen  atoms  are  dis- 
placed. Thus,  there  are  two  sodium  sulphites — 

(1)  Normal  sodium  sulphite,  Na2SO3. 

(2)  Hydrogen    sodium   sulphite    (or    sodium    bi-sulphite), 
HNaSO3. 

Salts,  like  this  hydrogen  sodium  sulphite,  in  which  only 
a  part  of  the  hydrogen  originally  present  in  the  acid  has  been 
exchanged,  are  sometimes  called  acid  salts.  It  does  not 
follow,  however,  that  they  have  an  acid  reaction  towards 
litmus,  although  in  many  cases  they  do  happen  to  possess 
this  property.  It  must  be  remembered  that  an  acid  salt 
simply  means  a  salt  in  which  the  whole  of  the  hydrogen  of 
the  acid  has  not  been  displaced  by  the  metal. 

Sulphites  are  all  decomposed  by  stronger  acids,  such  as 
hydrochloric  or  sulphuric. 

The  action  of  sulphuric  acid  is  shown  by  the  equation 
above.  With  hydrochloric  acid  the  only  difference  is  that 
a  chloride  of  the  metal  is  formed,  thus — • 

K2S03  +  2HC1  -  2KCI  +  H2O  +  SO2. 
Note  that  when  lead  dioxide  and   sodium  dioxide  were 
burnt  in  sulphur  dioxide,  the  sulphates  and  not  the  sulphites 
of  the  metals  were  produced. 


CHAPTER   XXIX. 

SULPHUR   TRIOXIDE — SULPHURIC    ACID — SULPHATES. 

Sulphur  Trioxide,  SO3. — When  sulphur  burns  in  oxygen, 
however  much  oxygen  there  may  be,  the  compound  produced 
is  always  sulphur  dioxide,  and  only  a  minute  trace  of  the 
trioxide  is  formed  at  the  same  time.  But  if  a  mixture  of 
sulphur  dioxide  and  oxygen  is  passed  through  a  heated  tube 
containing  very  finely  divided  platinum,  then  the  sulphur 
dioxide  combines  with  the  oxygen  and  gives  sulphur  trioxide. 
The  way  in  which  the  heated  platinum  causes  these  two  gases 
to  unite  together  is  not  clearly  known,  but  the  platinum  itself 
is  not  altered. 

Experiment  245. — Dip  some  asbestos  fibres  into  a  solution  of 
platinum  chloride,  and  then  hold  them  by  means  of  a  small  pair 
of  tongs  in  a  bunsen  flame  until  they  are  quite  hot.  The  platinum 
chloride  first  dries  and  then  decomposes,  leaving  the  asbestos 
coated  over  with  very  finely  divided  platinum. 

Now  pack  a  quantity  of  this  "  platinized  asbestos  "  into  a  bulb 
tube,  which  is  supported  as  shown  in  Fig.  38.  Remove  the  bottle 
w  from  the  sulphur  dioxide  apparatus  (Fig.  107),  and  replace  it  by 
a  similar  bottle  having  a  third  tube,  which  dips  into  the  acid. 
Connect  this  tube  with  a  supply  of  oxygen,  so  as  to  let  both  sulphur 
dioxide  and  oxygen  bubble  through  the  same  bottle  and  become 
mixed.  Notice  that  so  long  as  the  platinum  is  cold,  no  fumes 
of  sulphur  trioxide  escape  from  the  bulb  tube,  but  as  soon  as  it  is 
heated,  white  fumes  make  their  appearance.  If  these  fumes  are 
passed  through  a  U-shaped  tube,  kept  cold  by  being  placed  in  a 
freezing  mixture  (powdered  ice  and  salt),  white  silky -looking  crystals 
will  condense  in  the  cold  tube.  This  is  the  sulphur  trioxide. 

Sulphur  trioxide  has  a  most  powerful  affinity  for  water. 
If  exposed  to  the  air,  it  soon  takes  enough  moisture  from  the 


256  Sulphuric  Acid. 

air  to  convert  itself  into  sulphuric  acid.  If  the  crystals  of 
sulphur  trioxide  are  dropped  into  water  they  combine  with 
great  energy,  making  a  hissing  sound  like  a  red  hot  iron  going 
into  water.  If  placed  upon  the  skin  it  produces  painful  burns. 

Sulphuric  Acid,  H2SQ4,  is  the  most  important  of  all  the 
sulphur  compounds,  and  its  manufacture  is  carried  on  on  an 
enormous  scale. 

The  frequency  with  which  we  have  used  this  powerful 
acid  substance  in  order  to  bring  about  chemical  decompositions 
cannot  fail  to  have  been  noticed.  It  was  used  in  the  prepara- 
tion of  hydrogen,  chlorine,  hydrochloric  acid,  nitric  acid, 
carbon  monoxide,  and  sulphur  dioxide. 

Modes  of  Formation. — Sulphuric  acid  is  produced  when 
sulphur  trioxide  is  dissolved  in  water. 

S03  +  H20  =  H2S04. 

It  is  also  produced  slowly,  by  the  gradual  absorption  of 
oxygen  by  a  solution  of  sulphurous  acid  (see  Sulphurous 
Acid). 

The  process  by  which  it  is  prepared  on  a  manufacturing 
scale,  consists  in  making  sulphur  dioxide  combine  with  oxygen 
(from  the  air)  in  the  presence  of  water  (steam). 

S02  +  O  +  H20  =  H2S04. 

We  have  seen  already  that  sulphur  dioxide  and  oxygen 
do  not  combine  very  readily,  but  require  help  in  order  to  make 
them  unite.  Therefore,  if  sulphur  dioxide,  oxygen,  and  steam 
were  simply  mixed  together,  they  would  be  a  very  long  time 
in  uniting  to  form  sulphuric  acid. 

The  compound  employed  to  cause  the  sulphur  dioxide  to 
combine  with  oxygen  is  nitric  oxide,  NO. 

We  have  learnt  already  (p.  199)  that  when  nitric  oxide 
comes  in  contact  with  the  air,  it  unites  with  another  atom  of 
oxygen,  and  forms  nitrogen  peroxide,  NO^  a  reddish  gas. 
Now  nitrogen  peroxide  easily  gives  up  this  extra  atom  of 
oxygen  to  sulphur  dioxide  in  the  presence  of  steam,  and  goes 
back  again  to  nitric  oxide.  Therefore,  when  steam,  sulphur 


Manufacture  of  Sulphuric  Acid.  257 

dioxide,  and  nitrogen  peroxide  are  mixed,  the  following  change 
takes  place — 

H20  +  S02  +  N02  =  H2S04  +  NO. 

The  nitric  oxide  that  is  thus  formed  instantly  seizes 
another  atom  of  oxygen  from  the  air,  again  forming  nitrogen 
peroxide,  NO2, 

NO  +  O  =  NO2 

and  this  again  hands  on  this  extra  oxygen  to  another  portion 
of  sulphur  dioxide. 

The  nitric  oxide  (NO)  is,  therefore,  a  sort  of  "  middleman," 
who  takes  oxygen  from  the  air  and  passes  it  on  to  the  sulphur 
dioxide ;  and  the  same  quantity  of  nitric  oxide  can  keep  on 
doing  this,  and  can  convert  an  unlimited  amount  of  sulphur 
dioxide  into  sulphuric  acid. 

The  Manufacture  of  Sulphuric  Acid  is  carried  on  in 
enormous  rooms  made  of  sheet  lead.  These  great  rooms  are 
called  leaden  chambers,  and  they  are  often  of  such  a  size  that 
250  people  could  sit  down  to  dine  in  one.  Generally  several 
are  placed  in  a  row. 

Into  these  chambers  are  sent  sulphur  dioxide,  air,  nitrogen 
peroxide,  and  steam. 

(1)  The   sulphur  dioxide  is  produced  either  by  burning 
sulphur,  or  roasting  iron  pyrites,  in   special  furnaces  called 
sulphur  burners   or  pyrites  burners.     In   either   case   sulphur 
dioxide  is  formed,   which,   with   the   excess    of  air   passing 
through  the  furnace,  is  drawn  into  the  "  chambers." 

S  +  O2  =  SO2. 

Sometimes  the  sulphur  dioxide  is  obtained  by  burning 
sulphuretted  hydrogen. 

(2)  The  nitrogen  peroxide  is  produced  by  placing  inside  one 
of  the  sulphur  burners  a  pot  containing  a  little  Chili  saltpetre 
and  sulphuric   acid.     This  mixture  when  heated  gives  nitric 
acid  (p.  189),  and  the  fumes  of  nitric  acid  coming  in  contact 
with   the  sulphur  dioxide  are  decomposed,  yielding  nitrogen 
peroxide,  which  passes  on  into  the  chambers. 

2  HNO3  -f  SO2  =  H2SO4  4-  2  NO,. 

s 


258  Sulphuric  Acid. 

(3)  The  air  that  is  admitted  into  the  chambers  is  what  is 
allowed  to  pass  through  the  pyrites  burners,  and  its  amount  is 
regulated. 

(4)  The  steam  is  blown  into  the  chambers  in  jets  from 
steam  boilers. 

When  these  gases  mix  in  the  chambers,  the  chief  reactions 
which  go  on  are  represented  by  the  equations  given  above, 
and  sulphuric  acid  (moderately  strong)  collects  on  the  floors 
and  is  drawn  off.  Care  is  taken  to  adjust  the  proportion  of 
the  various  gases. 

As  the  oxides  of  nitrogen  go  on  transferring  oxygen  from 
the  air  to  the  sulphur  dioxide  over  and  over  again,  it  is  only 
necessary  to  add  a  very  small  amount  of  the  nitrogen  peroxide 
from  time  to  time,  to  make  up  for  the  slight  loss  of  this  gas 
which  always  takes  place.1 

The  acid  which  is  drawn  from  the  chambers  (chamber  acid} 
is  boiled  down,  either  in  glass  or  platinum  vessels,  so  as  to 
drive  off  the  water,  and  so  get  strong  sulphuric  acid. 

Properties. — Pure  sulphuric  acid  is  a  heavy,  colourless, 
oily  liquid  (hence  the  common  name  oil  of  vitriol}.  It  is  a 
powerfully  corrosive  substance,  and  if  spilt  upon  the  skin 
produces  bad  burns.  Therefore  some  care  must  be  taken  in 
handling  this  acid. 

It  has  a  strong  affinity  for  water,  and  if  mixed  with  water 
the  mixture  gets  nearly  boiling  hot.  On  account  of  its  power 
of  combining  witli  water,  it  is  constantly  used  for  withdrawing 
water  vapour  from  gases.  Thus,  when  we  require  to  dry  a  gas, 
that  is,  to  remove  the  vapour  of  water  from  it,  we  bubble  the 
gas  through  sulphuric  acid,  provided  it  is  a  gas  which  has  no 
action  on  the  acid.  If  this  acid  is  exposed  to  the  air,  or  is  not 
kept  in  well-stoppered  bottles,  it  quickly  absorbs  water  vapour 
from  the  air,  and,  of  course,  by  so  doing  gets  more  and  more 
dilute. 

Its  affinity  for  water  is  so  great  that  it  decomposes  many 
compounds  containing  hydrogen  and  oxygen,  and  takes  these 
elements  away  from  the  compound  in  the  proportion  to  yield 

1  For  further  details  of  the  manufacturing  process,  see  "Newth's 
Inorganic  Chemistry." 


Sulphates.  259 

water.  Thus,  in  the  case  of  oxalic  acid  (p.  233),  C2H2O4. 
This  is  decomposed  by  sulphuric  acid,  which  in  its  eagerness 
for  water,  abstracts  from  the  oxalic  acid  the  elements  which 
yield  water,  H2  and  O,  leaving  just  enough  oxygen  for  the 
carbon  to  form  carbon  monoxide  and  carbon  dioxide. 

Again  in  the  case  of  sugar  (p.  212),  C^H^Ou.  When 
sulphuric  acid  acts  on  this,  it  abstracts  all  the  hydrogen  and 
oxygen  (which  are  present  in  exactly  the  proportion  to  give 
nH2O)  and  leaves  the  carbon  in  the  free  or  uncombined 
state.  Its  power  of  charring  organic  matter  may  be  shown 
by  the  following  experiment. 

Experiment  246. — Take  some  very  dilute  sulphuric  acid,  and 
with  the  finger  write  a  word  on  a  piece  of  paper.  Now  gently  dry 
the  paper  by  holding  it  at  some  distance  above  a  gas  flame.  Notice 
that  the  paper  is  charred  where  the  letters  were  drawn  upon  it. 

Sulphates. — Just  as  carbonic  acid  and  sulphurous  acid 
form  two  classes  of  salts,  so,  for  the  same  reason,  there  are  two 
classes  of  sulphates.  The  reason  being  that  sulphuric  acid, 
like  these  others,  contains  two  atoms  of  hydrogen  which  can 
be  replaced  by  metals. 

Thus,  by  replacing  the  hydrogen  atoms  by  potassium  we 
get  either  normal  potassium  sulphate  (or  potassium  sulphate), 
K2SO4 ;  or  hydrogen  potassium  sulphate  (potassium  bi-sulphate), 
HKSO4. 

Certain  of  the  sulphates  were  known  to  the  very  early 
chemists,  and  were  called  vitriols  (because  they  had  rather  a 
vitreous  or  glassy  appearance),  such  as  blue  vitriol  (copper 
sulphate),  green  vitriol  (iron  sulphate),  white  vitriol  (zinc 
sulphate).  The  name  "  oil  of  vitriol "  is  derived  from  the  fact 
that  the  acid  was  formerly  obtained  by  distilling  green  vitriol. 

The  sulphates,  like  all  other  salts,  are  formed  when  the 
acid  is  neutralized  with  a  metallic  hydroxide. 

2KHO  +  H2S04  =  K2S04  +  2H20. 

They  are  also  produced  by  the  action  of  the  acid  on  oxides 
and  carbonates — 

ZnO  +  H2SO4  =  ZnSO4  +  H2O; 
Na2CO3  -f  H2SO4  =  Na2SO4  +  CO2  -f  HaO. 


260  Sulphuric  Acid. 

Being  such  a  powerful  acid,  sulphuric  acid  is  capable  of 
taking  metals  away  from  the  salts  of  almost  any  other  acid ; 
for  example,  from  sodium  chloride  or  potassium  nitrate  it 
takes  the  sodium  or  potassium,  and  gives  its  hydrogen  in 
exchange.  Thus  — 

2NaCl  +  H2SO4  =  Na2SO4  +  2HC1  (p.  127); 
2KNO3  +  H2SO4  =  K2SO4    +  2HNO3  (p.  189). 

Test  for  Sulphates. — (i)  Sulphates  which  are  soluble 
in  water  are  recognized  and  distinguished  by  giving  a  white 
precipitate  with  barium  chloride,  consisting  of  barium  sulphate. 
This  white  precipitate  is  insoluble  in  either  hydrochloric  or  nitric 
acid. 

Experiment  247. — Dissolve  in  three  separate  test-tubes  a  small 
particle  of  (i)  potassium  sulphate,  (2)  sodium  sulphite,  and  (3) 
sodium  carbonate.  Add  to  each  a  few  drops  of  barium  chloride. 
Notice  that  a  very  similar  precipitate  is  formed  in  each  case,  but 
in  reality  they  are  totally  different — one  is  barium  sulphate,  the 
next  is  barium  sulphite,  and  the  third  is  barium  carbonate.  To 
each  add  a  few  drops  of  strong  hydrochloric  acid.  Notice  that  the 
barium  sulphate  is  unaffected  ;  the  barium  carbonate  quickly  dis- 
solves with  effervescence,  giving  off  carbon  dioxide  ;  while  the 
barium  sulphite  also  dissolves  with  effervescence,  and  gives  off 
sulphur  dioxide  (which  can  be  detected  by  the  smell).  [Probably 
in  this  case  the  precipitate  will  not  wholly  dissolve,  because  the 
sodium  sulphite  originally  used  is  likely  to  contain  a  little  sodium 
sulphate  mixed  with  it,  so  that  the  precipitate  obtained  when  barium 
chloride  was  added,  consists  partly  of  barium  sulphite  (which  will 
dissolve  in  the  acid)  and  barium  sulphate  which  will  not  dissolve.] 
The  reactions  in  this  experiment  are  the  following. 

(«)  When  barium  chloride  is  added  to  the  three  solutions — 

K2SO4    +  BaCl2  =  2KC1   +  BaSO4  white  precipitate 
Na2SO3  +  BaCl2  =  2NaCl  +  BaSO3      „ 
Na2C03  +  BaCl2  =  2NaCl  +  BaCO3      „ 

(b)  The  action  of  hydrochloric  acid  on  the  three  precipitates— 

BaSO4  +  2HC1  no  action. 

BaSO3  +  2HCI  =  BaCl2  +  H2O  +  SO2 

BaCO3  +  2HC1  =  BaCl2  +  H2O  +  CO2. 


Tests  for  Sulphates. 


261 


(2)  Sulphates  which  are  not  soluble  in  water  are  tested  for 
in  a  different  way. 

Experiment  248. — Take  a  pinch  of  plaster  of  Paris  (calcium 
sulphate)  and  mix  it  with  about  three  times  as  much  powdered 
sodium  carbonate,  and  heat  the  mixture  on  a  little  piece  of  platinum 
foil,  bent  into  a  sort  of  spoon  (or  on  the  lid  of  a  platinum  crucible) 
until  it  has  completely  melted.  At  this  high  temperature  the  two 
compounds  make  a  mutual  exchange,  resulting  in  the  formation  of 
sodium  sulphate  and  calcium  carbonate. 

CaSO4  +  Na2CO3  =  Na2SO4  +  CaCO3. 

The  sodium  sulphate  is  soluble  in  water,  while  the  calcium  car- 
bonate is  not ;  therefore,  put  the  platinum  spoon  into  water  in  a 
test-tube  and  boil  it  for  a  minute 
or  two,  and  then  filter  it.  The 
solution  contains  sodium  sul- 
phate, and  any  excess  of  sodium 
carbonate  which  may  have  been 
used.  Now  add  hydrochloric  acid 
until  the  solution  is  quite  acid, 
and  then  add  barium  chloride. 
The  white  precipitate  of  barium 
sulphate  is  at  once  formed.  [In- 
stead of  fusing  the  insoluble  sul- 
phate with  sodium  carbonate,  we 
may  mix  it  with  a  strong  solution 
of  sodium  carbonate  and  boil  it 
for  some  time,  when  there  will  be 
enough  of  the  calcium  sulphate 
decomposed  and  sodium  sulphate 
produced  to  give  the  test  with 
barium  chloride.] 

Experiment  249. — Heat  a  little  heap  of  a  mixture  of  plaster  of 
Paris  and  sodium  carbonate  on  a  piece  of  charcoal  by  means  of 
a  small  blow-pipe  flame,  as  shown  in  Fig.  108.  [Select  a  good 
piece  of  charcoal,  which  is  not  full  of  cracks,  and  scoop  a  small 
hollow  upon  it  to  hold  the  substance  which  is  being  heated.]  The 
result  of  this  operation,  is  that  the  sodium  of  the  sodium  carbonate 
combines  with  the  sulphur  of  the  calcium  sulphate  to  give  sodium 
sulphide.  When  it  is  cold,  place  the  fused  residue  on  a  silver  coin 
and  touch  it  with  a  drop  of  water.  The  sodium  sulphide  at  once 
acts  on  the  silver  and  causes  a  black  stain  upon  it. 


FIG.  108. 


262  Sulphuric  Add. 


EPITOME. 

The  two  commonest  oxides  of  sulphur  are  sulphur  dioxide,  SO2, 
and  sulphur  trioxide,  SO3.  Sulphur  dioxide  is  a  colourless  choking 
gas,  obtained  when  sulphur  burns  either  in  air  or  in  oxygen.  It  is 
prepared  by  heating  sulphuric  acid  and  copper.  It  is  more  than 
twice  as  heavy  as  air,  and  therefore  can  be  collected  by  downward 
displacement. 

It  is  soluble  in  water,  and  therefore  cannot  be  collected  at  the 
pneumatic  trough. 

Sulphur  dioxide  does  not  burn,  nor  support  the  combustion  of 
ordinary  burning  bodies.  Lead  dioxide  and  sodium  dioxide  take 
fire  in  the  gas  and  produce  sulphates  of  the  metals. 

Sulphur  dioxide  bleaches,  but  not  always  permanently  ;  the 
colour  often  being  restored  either  by  stronger  acids  or  by  alkalies. 

The  gas  is  used  for  disinfecting  purposes  ;  but  its  chief  use  is 
in  the  manufacture  of  sulphuric  acid. 

Sulphur  dioxide  is  easily  condensed  to  a  liquid  by  cooling  it. 
The  liquefied  gas  is  colourless,  and  boils  at  —8°. 

The  solution  of  sulphur  dioxide  in  water  is  acid,  and  contains 
sulphurous  acid,  H2S03.  This  acid  has  never  been  obtained  except 
as  a  solution  in  water.  When  the  solution  is  boiled,  the  acid  is 
decomposed  and  sulphur  dioxide  escapes. 

The  salts  of  sulphurous  acid  are  called  sulphites.  The  acid  is 
dibasic,  and  therefore  forms  two  classes  of  salts  ;  those  in  which 
all  the  hydrogen  has  been  replaced  by  metals,  and  those  in  which 
only  half  the  hydrogen  is  so  replaced.  Thus  :  Na2SO3,  di-sodium 
sulphite,  or  normal  sodium  sulphite  ;  and  HNaSO3,  hydrogen 
sodium  sulphite  (sometimes  called  add  sodium  sulphite,  or  sodium 
bisulphite) . 

The  sulphites  are  decomposed  by  hydrochloric  or  sulphuric  acid, 
with  the  evolution  of  sulphur  dioxide. 

Sulphur  trioxide  is  a  white  solid,  forming  silky  crystals.  It  is 
produced  when  sulphur  dioxide  and  oxygen  are  passed  over  heated 
spongy  platinum.  This  oxide  has  a  powerful  affinity  for  water,  with 
which  it  combines  to  form  sulphuric  acid. 

Sulphuric  acid  (or  oil  of  vitriol)  is  a  colourless  oily  liquid, 
strongly  corrosive,  and  a  powerful  acid.  Its  manufacture  is  the 
most  important  of  all  chemical  industries.  It  is  made  by  oxidizing 
sulphur  dioxide  by  means  of  nitrogen  peroxide  in  the  presence  of 
steam.  Sulphur  dioxide  combines  with  atmospheric  oxygen,  only 
with  extreme  slowness  ;  but,  by  means  of  nitric  oxide,  oxygen  is 


Sulphuric  Acid.  263 

taken  from  the  air  and  handed  on  to  the  sulphur  dioxide.  Nitric 
oxide  unites  with  oxygen  of  the  air,  forming  nitrogen  peroxide,  and 
this  gives  oxygen  to  the  sulphur  dioxide,  and  is  again  reduced  to 
nitric  oxide. 

The  operation  is  carried  on  in  enormous  chambers  built  of 
lead. 

Sulphuric  acid  combines  with  water  with  the  production  of  great 
heat.  On  account  of  its  eagerness  to  unite  with  water  it  is  used  for 
drying  gases.  It  also  decomposes  many  organic  compounds  con- 
taining oxygen  and  hydrogen,  withdrawing  these  two  elements  in 
the  proportion  required  to  form  water. 

Sulphuric  acid  is  dibasic,  and  forms  two  classes  of  salts,  accord- 
ing as  to  whether  all,  or  only  half,  the  hydrogen  of  the  acid  is 
replaced  by  metals. 


CHAPTER   XXX. 

SOME   COMMON   CARBON    COMPOUNDS. 

CARBON  forms  such  an  enormous  number  of  compounds  that 
a  simple  list  of  they:  names  alone  would  more  than  fill  this 
little  book.  The  study  of  these  compounds  is  generally  called 
"  organic  chemistry,"  because,  in  the  early  days  of  chemistry, 
it  was  thought  that  these  compounds  could  only  be  produced 
as  the  result  of  living  organized  bodies,  like  animals  and 
plants. 

In  order  to  study  this  vast  host  of  compounds,  they  are 
divided  and  subdivided  into  classes  and  families  much  in  the 
same  way  as  animals  or  plants  are  classified. 

Four  very  important  classes  are  :— 

1.  Hydrocarbons. 

2.  Acids. 

3.  Alcohols. 

4.  Carbohydrates. 

1.  Hydrocarbons. — These,  as  the  name  implies,  are  com- 
pounds of  carbon  with  hydrogen  only.  Important  amongst 
these  are  the  following  : — 

Marsh  Gas,  CH4. — Found  in  coal  mines,  and  called 
fire-damp.  Also  in  marshy  places  where  vegetable  matter  is 
rotting.  In  the  laboratory  it  is  made  by  strongly  heating  a 
mixture  of  sodium  acetate  and  caustic  soda,  when  sodium 
carbonate  is  left  behind. 

NaC2H3O2  +  NaHO  =  Na2CO3  +  CH4. 

Marsh  gas  is  colourless,  and  has  no  smell.  It  burns  easily, 
but  gives  very  little  light,  although  rather  more  than  hydrogen. 


Common  Carbon  Compounds.  265 

If  mixed  with  air,  or  with  pure  oxygen,  and  fired,  the  mixture 
explodes.  This  is  the  cause  of  coal-mine  explosions.  Water 
and  carbon  dioxide  are  produced,  and  the  latter  gas  is  called 
choke-damp  by  the  miners. 

CH4+  3O  =  2H2O  +  COo. 

Ethylene,  C2K4  (olefiant  gas),  is  obtained  from 
common  alcohol  (spirits  of  wine)  by  heating  it  with  sulphuric 
acid  (see  Exp.  184). 

C2H60  -  H20  =  C2H4 

Ethylene  is  a  colourless  gas  with  a  faint  pleasant  smell.  It 
burns  easily,  with  a  very  bright  flame,  producing  carbon  dioxide 
and  water,  the  same  products  as  are  formed  when  any  hydro- 
carbon burns. 

Acetylene,  CSH2,  is  formed  whenever  coal  gas  burns 
without  a  sufficient  supply  of  air.  Thus,  when  a  Bunsen  lamp 
gets  alight  down  at  the  little  jet  at  the  base  of  the  chimney, 
some  acetylene  is  produced.  This  it  is  which  causes  the  bad 
smell  we  notice  when  the  lamp  so  burns. 

Acetylene  is  best  prepared  in  the  laboratory  by  acting  on 
calcium  carbide  with  water.  A  little  of  the  carbide,  in  small 
lumps,  is  put  into  a  test-tube,  and  a  few  drops  of  water  added 
The  gas  is  at  once  evolved,  and  can  be  lighted  at  the  mouth  of 
the  tube. 

CaC2  +  2H2O  -  CaH2O2  +  C2H2. 

Acetylene  burns  with  a  very  bright  and  smoky  flame. 

Marsh  gas,  Ethylene,  and  Acetylene  are  all  present  in 
ordinary  coal  gas. 

Other  important  hydrocarbons  are  the  various  mineral  oils 
(paraffin  oils)  used  for  illuminating  purposes.  Paraffin  wax, 
used  for  candles.  Turpentine. 

2.  The  Acids. — This  is  a  large  and  important  class, 
divided  into  many  families.  Amongst  the  most  important  of 
these  is  the  one  known  as  the  acetic,  or  fatty  series  of  acids. 
These  include  acids  present  in  many  fats. 

Formic  Acid  (CH2O2)  is  found  in  ants,  and  in  the  hairs 


266  Common  Carbon  Compounds. 

of  stinging-nettles.  When  we  are  stung  by  ants  or  nettles,  it 
is  because  a  minute  drop  of  this  formic  acid  has  been  injected 
into  the  skin.  If  a  piece  of  litmus  paper  is  placed  upon  an 
ants'  nest  which  has  been  just  disturbed,  it  is  instantly  reddened 
by  a  shower  of  formic  acid  being  squirted  against  it  by  the 
irritated  ants.  Formic  acid  is  decomposed  by  strong  sulphuric 
acid  into  water  and  carbon  monoxide. 

Acetic  Acid  i(C2H4O2)  is  the  acid  of  vinegar.  When 
beer  goes  sour  it  is  owing  to  the  formation  of  acetic  acid,  and 
this  change  is  brought  about  by  the  agency  of  a  living  organism 
familiarly  known  as  mother  of  vinegar.  Acetic  acid  is  also 
produced  when  wood  is  destructively  distilled,  the  material  so 
obtained  being  called  pyroligneous  arid,  that  is,  the  fire-wood 
acid. 

Pure  acetic  acid  is  liquid  at  the  temperature  of  a  warm 
room ;  but  in  winter  it  freezes  to  an  ice-like  solid.  It  is  on 
this  account  called  "  glacial "  acetic  acid.  This  strong  acid 
stings  the  skin  as  formic  acid  does.  The  salts  are  called 
acetates ;  one  of  the  commonest  is  lead  acetate,  familiarly 
known  as  sugar  of  lead. 

Butyric  Acid  is  the  name  of  the  acid  which  is  present  in 
rancid  butter,  and  which  gives  to  it  the  disagreeable  smell. 

Palmitic  Acid  and  Stearic  Acid  are  important  con- 
stituents of  most  solid  animal  fats,  such  as  beef  and  mutton 
suet,  butter,  and  human  fat.  Palmitic  acid  is  also  one  of 
the  chief  constituents  of  palm  oil  (hence  its  name).  These 
two  acids  are  extensively  used  for  making  candles. 

3.  Alcohols.  This  class  contains  a  number  of  very 
important  compounds.  The  following  may  be  taken  as 
examples. 

Methyl  Alcohol  (wood  spirit},  CH4O  or  CH3HO,  is 
present  in  the  watery  liquid  obtained  when  wood  is  distilled. 
When  it  is  extracted  from  this  liquid  and  purified,  it  is  a 
colourless  liquid,  which  burns  with  a  flame  without  light,  and 
without  any  smoke. 

Ethyl  Alcohol,  C2H6O  or  C2H5HO,  is  the  familiar 
spirits  of  wine.  It  is  the  best  known  of  all  the  alcohols,  and 
is  therefore  called  simply  "  alcohol." 


Methylated  Spirit.  267 

It  is  obtained  by  the  fermentation  of  ordinary  sugar  or  of 
grape  sugar  (glucose)  by  means  of  yeast.  The  yeast  organism 
transforms  the  sugar  into  alcohol  and  carbon  dioxide  (see 
Exp.  195). 

C6H1206  =  2C2H60  +  2C02. 

Glucose.  Alcohol. 

The  liquid  obtained  contains  only  a  small  proportion  of 
alcohol  mixed  with  a  large  quantity  of  water.  It  has,  there- 
fore, to  be  distilled,  or  "rectified,"  in  order  to  separate  the 
alcohol  from  the  water. 

Ethyl  alcohol  is  present  in  all  fermented  liquors,  and  it  is 
the  presence  of  this  compound  which  gives  them  their  intoxi- 
cating properties.  Ordinary  beer  contains  from  3  to  6  per 
cent,  of  alcohol.  [If  some  beer  is  boiled  in  a  flask  with  a 
long  upright  tube  fastened  into  the  neck  with  a  cork,  the 
alcohol  which  first  boils  off  can  actually  be  lighted  as  it 
escapes  from  the  tube.]  Light  wines,  such  as  claret,  etc., 
contain  from  8  to  14  per  cent.;  port  and  sherry  15  to  25  per 
cent,  while  brandy  and  other  "  spirits  "  contain  from  50  to  60 
per  cent  of  alcohol. 

Alcohol  is  capable  of  dissolving  things  like  resins,  gums, 
oils,  etc.,  and  is  therefore  a  most  useful  substance  for  the 
manufacture  of  varnishes,  and  for  other  purposes. 

Methylated  Spirit. — On  account  of  the  high  duty  upon 
"alcohol,"  it  is  too  expensive  for  most  of  the  manufacturing 
and  chemical  purposes  for  which  alcohol  is  required.  There- 
fore a  mixture  consisting  of  90  per  cent.  "  spirits  of  wine " 
and  10  per  cent.  "  wood  spirit "  (impure  methyl  alcohol)  is 
used  instead  of  pure  "  alcohol."  This  mixture  called  methylated 
spirit  is  quite  unfit  for  drinking,  and,  being  duty  free,  is  quite 
cheap. 

Phenol  or  Phenyl  Alcohol,  C6H5HO.— This  substance 
is  familiarly  known  by  the  name  carbolic  arid,  although  in 
reality  it  is  not  an  acid,  but  belongs  to  the  class  of  alcohols. 
It  is  formed  when  coal  is  distilled  for  making  coal-gas,  and 
is  extracted  from  the  coal-tar  oil. 

Pure  phenol  forms  long  needle-shaped  crystals,  which 
melt  a  little  above  the  temperature  of  the  hand  (42°).  It  has 


268  Common  Carbon  Compounds. 

a  sharp  burning  taste  and  is  poisonous.  It  is  a  powerful 
disinfectant  and  antiseptic,  and  is  largely  used  in  surgery. 

The  common  "  carbolic  acid  powders "  sold  in  tins  for 
disinfecting  purposes,  consist  of  some  powder  such  as  lime 
or  gypsum,  impregnated  with  about  10  per  cent,  of  very 
crude  carbolic  acid. 

Glycerin,  C3H5(HOJ3.— This  familiar  substance  is  also 
an  alcohol.  Chemists  call  it  glycerol.  It  is  an  important 
constituent  of  fats.  Fats  are  compounds  of  glycerin  with  such 
acids  as  palmitic  and  stearic.  Mutton  suet,  for  instance, 
consists  chiefly  of  a  compound  of  stearic  acid  and  glycerin. 
This  compound  is  called  stearin.  Palmitin  is  the  name  of  the 
compound  of  palmitic  acid  with  glycerin.  Glycerin  is  obtained 
from  fats  by  heating  them  in  very  hot  steam,  or  by  boiling 
them  with  a  caustic  alkali  such  as  sodium  hydroxide.  The 
alkali  combines  with  the  stearic  or  palmitic  acid  forming 
a  salt  (soap),  which  separates  out  when  the  liquid  cools  ;  while 
the  glycerin  which  is  set  free  remains  dissolved  in  the  watery 
liquid. 

Glycerin  is  a  thick  syrupy  liquid,  with  a  very  sweet  taste, 
almost  like  sugar  syrup. 

Soap. — When  fats  are  boiled  with  caustic  alkalies,  the  fat 
(which  is  a  compound  of  fatty  acids  with  glycerin)  is  decom- 
posed ;  glycerin  is  set  free  and  the  fatty  acid  combines  with 
the  alkali.  This  process  is  called  saponification ;  and  the 
compound  of  the  fatty  acids  (chiefly  stearic  and  palmitic  acids) 
with  the  alkali,  is  known  as  a  soap. 

When  caustic  potash  is  the  alkali  used  soft  soap  is  pro- 
duced ;  while  if  caustic  soda  is  employed  the  soap  is  harder, 
as  in  the  ordinary  forms  of  soap  used  for  washing  purposes. 

In  some  out  of  the  way  parts  of  the  world,  people  often 
make  a  crude  kind  of  soap  by  collecting  the  ashes  of  burnt 
wood  (these  contain  potash,  the  name  potash  simply  mean- 
ing pot-ashes),  mixing  them  with  water,  and  boiling  the  liquid 
so  obtained  with  mutton  or  beef  fat 

The  Action  of  Hard  Waters  upon  Soap.— The 
property  of  hardness  in  water  is  chiefly  due  to  the  presence 
of  either  carbonate  of  lime  or  sulphate  of  lime  (see  p.  81). 


Soap.  269 

When  water  containing  these  salts  in  solution  comes  in  contact 
with  a  solution  of  soap  (say  sodium  stearate),  a  chemical  action 
takes  place,  thus  — 

Sodium  stearate  4-  calcium  carbonate  =  calcium  stearate  -f 

(Insoluble.) 

sodium  carbonate. 

The  calcium  stearate,  being  insoluble,  separates  out  as  a 
greasy  scum,  which  is  always  seen  when  hard  water  is  used 
for  washing.  So  long  as  any  of  the  lime  salts  remain  in  the 
water,  the  soap  is  used  up  in  bringing  about  this  double 
decomposition,  and,  therefore,  is  wasted  so  far  as  its  powers 
of  cleansing  are  concerned ;  for  we  cannot  wash  with  soap 
until  the  water  has  become  softened  by  the  removal  or 
destruction  of  the  hardening  lime  salts.  As  soon  as  ever 
all  the  calcium  carbonate  or  sulphate  has  been  decomposed 
by  the  soap,  then,  and  not  till  then,  will  the  soap  give  a  lather  ; 
and  not  until 'a  lather  can  be  produced  is  the  soap  of  any 
service  for  washing. 

Experiment  250. — Take  two  good-sized  stoppered  bottles,  and 
about  half  fill  one  with  some  hard  water  (common  tap-water  will 
generally  do,  but  if  this  should  happen  to  be  a  soft  water,  a  sample 
of  hard  water  can  readily  be  made  by  adding  a  little  lime-water, 
and  then  bubbling  carbon  dioxide  through  it).  At  first,  calcium 
•  carbonate  CaCO3  is  precipitated,  but  presently  this  dissolves  in  the 
carbonic  acid,  forming  the  bi-carbonate  of  calcium  H2Ca(CO3)2,  or 
CaCO3,H2CO3.  Into  the  other  bottle  put  an  equal  quantity  of 
distilled  water,  or  rain  water.  Now  make  a  solution  of  soap  by 
shaking  up  a  few  thin  shavings  of  soap  with  a  little  water  in  a 
bottle.  Add  a  small  quantity  of  this  soap  solution  at  a  time  to  the 
soft  water,  until  on  shaking  the  bottle  a  lather  is  raised  which  does 
not  disappear  again.  Now  add  the  soap  solution  to  the  hard 
water.  Note  that  at  first  no  lather  is  produced,  although  much 
more  of  the  soap  is  added.  Also  observe  that  the  solution  turns 
muddy,  and  a  scum  is  produced.  This  is  the  calcium  stearate 
being  precipitated ;  continue  adding  the  soap  until  at  last  a  per- 
manent lather  is  obtained.  If  a  little  more  hard  water  or  a  few 
drops  of  a  solution  of  any  lime  salt  be  now  added,  the  lather 
instantly  disappears. 


270  Common  Carbon  Compounds. 

4.  Carbohydrates. — This  family  of  carbon  compounds 
includes  the  sugars,  starches,  etc.  The  following  are  some 
common  examples  of  carbohydrates. 

Cane-sugar  (Saccharose),  C^H^On.  This  is  ordinary 
sugar,  and  is  obtained  chiefly  from  sugar-cane  and  from 
beet-root.  It  is  present  in  almost  all  sweet  fruits. 

When  cane-sugar  is  brought  into  contact  with  yeast,  it  is 
converted  into  dextrose  and  laevulose.  The  same  change 
takes  place  when  it  is  warmed  with  dilute  sulphuric  acid. 

C12H22On  +  H20  -  C6H1206-fC6H1206. 

Cane-sugar.  Dextrose.  Lsevulose. 

Grape-sugar  (Glucose  or  Dextrose],  C6H12O6,  is  obtained 
from  the  juice  of  sweet  grapes.  It  is  also  present  in  honey. 
It  is  not  so  sweet  as  cane-sugar,  and  is  less  easily  dissolved 
by  water.  Grape-sugar  is  distinguished  from  cane-sugar  by 
the  following  test :  — 

Experimental. — Take  a  crystal  of  copper  sulphate,  and  one 
about  the  same  size  of  tartaric  acid.  Dissolve  them  together  in  a 
little  water,  and  add  a  solution  of  caustic  potash  until  the  liquid 
is  strongly  alkaline.  Add  a  little  of  this  mixture  to  a  solution  of 
grape-sugar  in  a  test-tube,  and  warm  the  liquid.  A  red  precipitate 
of  cuprous  oxide  is  formed,  and  the  solution  loses  its  blue  colour. 
Do  the  same  with  a  solution  of  common  sugar  ;  this  does  not  give 
the  red  precipitate,  and  the  liquid  remains  blue. 

Starch,  C6H10O5,  occurs  in  many  parts  of  plants,  such  as 
the  seeds,  stems,  roots,  and  tubers,  and  it  may  be  obtained 
from  them  by  crushing  them  with  water,  and  separating  the 
bruised  fibre  or  pulp  by  means  of  a  sieve. 

Experiment  252. — Take  a  raw  potato  and  rub  it  down  on  a 
coarse  grater,  and  collect  all  the  pulp  on  a  piece  of  muslin.  Screw 
it  up  into  a  sort  of  bag,  and  squeeze  it  into  a  small  basin  of  water  ; 
dipping  it  once  or  twice  into  the  water  and  squeezing  again.  In 
this  way  the  starch  passes  through  the  muslin  into  the  water,  which 
is  thereby  made  milky.  If  this  is  allowed  to  stand,  the  clean  white 
starch  settles  to  the  bottom.  If  wheat  meal  or  flour  is  treated  in 
a  similar  manner,  starch  is  also  separated,  and  a  whitish  sticky 
mass  is  left.  This  is  called  gluten.  It  is  a  compound  containing 


Starch.  27 1 

nitrogen,  and  it  is  the  presence  of  this  substance  which  gives  to  the 
wheat  its  value  as  a  food.  Wheat  flour  contains  about  ^th  of 
its  weight  of  gluten. 

When  starch  is  boiled  with  dilute  sulphuric  acid,  or  is 
acted  on  by  the  ferment  present  in  germinating  barley 
(diastase),  it  is  converted  into  glucose  and  dextrin. 

3C6H1008+  H20  =  C6H1206  +  2C6H1005. 

Starch.  Glucose.  Dextrin. 

When  examined  through  a  microscope,  starch  is  found  to 
consist  of  minute  granules,  not  crystals.  The  size  and  shape 
of  these  vary  considerably.  Those  of  potato  starch  and 
arrowroot  are  much  larger  than  those  from  rice. 

Starch  gives  a  blue  colour  with  iodine  (see  p.  134). 

Starch  does  not  dissolve  in  cold  water,  but  with  boiling 
water  the  granules  swell  up  and  burst.  If  the  amount  of  hot 
water  is  very  large,  the  starch  disappears  and  seems  to  dis- 
solve, but  with  less  water  it  forms  a  gelatinous  or  pasty  mass. 

Dextrin. — When  starch  paste  is  boiled  with  a  little  dilute 
sulphuric  acid,  it  soon  becomes  thinner  and  thinner,  being 
changed  into  glucose  and  dextrin,  both  of  which  are  soluble. 
Dextrin  is  a  sticky  gummy  substance,  often  used  instead  of 
ordinary  gum  or  paste,  for  mounting  photographs  and  other 
similar  purposes.  It  is  sold  under  the  name  of  British  gum. 


CHAPTER   XXXI. 

SIMPLE   QUALITATIVE    ANALYSIS. 

THE  word  "analysis"  means  the  breaking  up  or  separation 
of  a  compound  into  its  components  or  elements  (see  p.  75). 
But  the  word  is  also  used  in  a  broader  sense,  and  is  applied 
to  any  processes  or  methods  by  which  the  chemist  is  able 
to  find  out  what  a  substance  is  composed  of,  in  order  to 
identify  that  substance.  For  example,  we  speak  of  "  micro- 
scopic analysis"  and  spectrum  analysis ;  these  are  not^ pro- 
cesses in  which  compounds  are  split  up  into  their  components, 
but  are  methods  which  enable  chemists  to  identify  different 
substances.  The  following  illustration  will  make  this  plain. 

Experiment  253. — Take  a  crystal  of  nitre  (potassium  nitrate), 
and  just  touch  the  edge  of  a  Bunsen  flame  with  it.  Notice  the  lilac 
or  violet  colour  which  it  imparts  to  the  flame.  This  colour  is 
characteristic  of  all  potassium  compounds  (compare  Exp.  44).  Place 
the  crystal  on  a  glass  plate,  dissolve  it  in  a  drop  of  warm  water, 
and  allow  the  solution  to  evaporate.  Examine  the  crystals  which 
are  formed  with  a  pocket  lens  or  microscope.  Notice  the  long 
prismatic  shaped  crystals,  characteristic  of  nitre.  Dissolve  a 
crystal  of  the  salt  in  water,  and  apply  the  test  for  a  nitrate,  as 
explained  on  p.  195. 

By  these  three  experiments  or  tests,  we  have  analysed  this 
substance,  and  identified  it  as  potassium  nitrate,  although  we 
have  not  separated  the  compound  into  its  constituent  elements. 

Separation  Usually  Necessary. — In  a  great  many 
instances,  the  tests  which  are  used  to  identify  a  substance 
are  interfered  with  if  certain  other  substances  are  present  at 
the  same  time.  In  all  such  cases  it  is  necessary  to  separate 


Reagents.  273 

the    substances    before    applying     the     special     test.       For 
example — 

Experiment  254. — Dissolve  a  small  crystal  of  potassium  chloride 
in  a  drop  or  two  of  water,  and  add  to  it  one  or  two  drops  of  a 
solution  of  platinum  chloride  (PtQ4).  A  precipitate  is  produced, 
consisting  of  tiny  yellow  crystals.  This  is  a  characteristic  test  for 
potassium  compounds. 

Now  treat  a  similar  quantity  of  ammonium  chloride  in  exactly 
the  same  way ;  notice  that  a  precisely  similar  looking  yellow 
crystalline  precipitate  is  formed. 

Therefore,  before  applying  this  test  for  the  detection  of  a 
potassium  salt,  it  is  absolutely  necessary  to  separate  it  from  any 
ammonium  salts. 

How  Separation  is  made. — The  method  most  fre- 
quently used  for  bringing  about  analytical  separation,  is  to 
cause  one  or  more  of  the  compounds  present  to  ^undergo  a 
double  decomposition  with  certain  chosen  reagents,  whereby 
the  metals  in  these  compounds  form  fresh  compounds  which 
are  insoluble,  in  water,  and  which  are,  therefore,  thrown  down 
as  precipitates.  An  example  of  this  method  is  given  on 
p.  20,  Exp.  25.  The  silver  nitrate  present  is  caused  to  enter 
into  double  decomposition  with  sodium  chloride.  This  results 
in  the  formation  of  silver  chloride,  which  is  precipitated  as 
an  insoluble  compound,  and  sodium  nitrate  remains  in  solution 
along  with  the  copper  salt.  It  must  be  noted  that  in  this 
separation  it  is  only  the  silver  from  the  silver  nitrate  that  is 
actually  separated  from  the  copper  compound,  for  the  other 
part  of  the  compound  is  left  in  solution  combined  with  sodium 
instead  of  with  silver.  By  this  process  we  have  not  separated 
silver  nitrate  from  copper  nitrate,  but  only  withdrawn  the 
silver  and  replaced  it  by  sodium. 

Reagents. — The  solutions  which  are  used  to  separate 
substances  in  this  way  are  called  reagents.  Sometimes  the 
same  reagent  will  form  insoluble  compounds  with  a  whole 
group  of  metals,  in  which  case  it  may  be  used  to  separate  an 
entire  family  of  metals  from  others  not  belonging  to  the  group. 
Such  reagents  are  called  group-reagents. 


274  Simple  Qualitative  Analysis. 

Groups. — For  convenience  the  metals  are  divided  into 
a  number  of  groups,  based  upon  their  behaviour  towards 
certain  chosen  group-reagents,  used  in  a  certain  order. 

GROUP  I.  or  HYDROCHLORIC  ACID  GROUP. — Metals  whose 
chlorides  are  precipitated  on  the  addition  of  hydrochloric  acid. 
Lead  1  (silver,  mercury). 

GROUP  II.  or  SULPHURETTED  HYDROGEN  GROUP.— 
Metals  whose  sulphides  are  precipitated  from  acid  solutions  by 
sulphuretted  hydrogen. 

Lead,2  Copper  (mercury,  bismuth,  cadmium,  tin,  arsenic, 
antimony). 

GROUP  Ilia,  or  AMMONIA  GROUP.  —  Metals  whose 
hydroxides  are  precipitated  by  ammonia,  in  the  presence 
of  ammonium  chloride. 

Iron  (chromium,  aluminium). 

GROUP  lllb.  or  AMMONIUM  SULPHIDE  GROUP. — Metals 
whose  sulphides  are  precipitated  by  ammonium  sulphide,  in 
the  presence  of  ammonia. 

Zinc  (manganese,  nickel,  cobalt). 

GROUP  IV.  or  AMMONIUM  CARBONATE  GROUP. — Metals 
whose  carbonates  are  precipitated  by  ammonium  carbonate, 
in  the  presence  of  ammonium  chloride. 

Calcium  (barium,  strontium). 

GROUP  V. — No  group  reagent. 

Potassium,  Ammonium  (sodium,  magnesium). 

Reactions  for  the  Metals.— In  order  that  we  may  be 
able  to  recognize  and  identify  a  metal  in  the  various  com- 
pounds it  produces  with  different  reagents,  it  is  necessary  to 
make  ourselves  quite  familiar  with  these  compounds.  In 
order  to  gain  this  knowledge,  the  following  reactions  or  tests 
should  be  carefully  made,  and  the  student  should  make  exact 
notes  of  all  he  does  and  observes.  If  any  experiment  he 
makes  seems  to  give  a  different  result  from  that  which  is 

1  Those  metals  printed  in  thick  type  are  the  only  ones  the  elementary 
student  will  be  concerned  with. 

2  The  reason  why  lead  is  in  both   groups  I.   and   II.  is   explained 
on  page  276. 


Reactions  for  Lead.  275 

indicated  in  the  book,  he  should  not  pass  it  over,  but  should 
repeat  it  more  carefully. 

Lead,  Pb.     (Group  I.) 

Use  lead  nitrate,  Pb(NO3)2.  Take  a  few  crystals  of  the 
salt ;  note  their  dry,  hard,  milk-white  appearance.  Dissolve 
a  little  in  water  in  a  test-tube.  Note  that  it  is  not  very  readily 
dissolved  in  cold  water,  but  more  quickly  if  the  water  is 
warmed.  Pour  some  of  this  solution  into  five  separate 
test-tubes. 

I.  To  one,  add  a  few  drops  of  dilute  hydrochloric  acid 
(group-reagent).      Note  the  white,  curdy   precipitate  of  lead 
chloride.      Shake   the   tube,   and,  when   the   precipitate   has 
settled,  add  more  of  the  reagent,  until  no  further  precipitation 
takes  place. 

Pb(NO3)2  4-  2HC1  =  PbCl2  +  2HN03. 

Now  heat  the  mixture,  and  observe  how  the  precipitate 
disappears.  Lead  chloride  is  soluble  in  hot  water.  This 
distinguishes  lead  from  the  other  metals  of  this  group.  Cool 
the  test-tube  again,  and  the  lead  chloride  is  again  precipitated  ; 
but  note  that  it  comes  down  in  the  form  of  shining  white 
needle-shaped  crystals. 

II.  Take  the  second  portion  of  the  lead  nitrate  solution, 
and  pass  sulphuretted  hydrogen  through  it  (see  p.  246),  or  add 
sulphuretted  hydrogen  water.      A  black   precipitate  of  lead 
sulphide  is  formed. 

Pb(N03)2  +  H2S  =  PbS  +  2HN03. 

Filter  the  liquid,  scrape  a  little  of  the  precipitate  off  the 
filter  into  a  test-tube,  add  a  few  drops  of  strong  nitric  acid, 
and  boil.  Notice  that  the  black  precipitate  turns  white. 
The  lead  sulphide  is  oxidized  by  the  nitric  acid  into  lead 
sulphate. 

III.  Treat  the  third  portion  as  the  first,   and   filter  the 
mixture.     Then   pass   sulphuretted   hydrogen  into   the  clear 
filtrate.     Note  a  black  precipitate,  as  in  II.     This  shows  that 
the    group    reagent    does    not  entirely  separate  lead    from 
group   II. ;  that  is  to  say,  lead  chloride   is  slightly  soluble 


276  Simple  Qualitative  Analysis. 

even  in  cold  water,  and,  therefore,  a   portion   of  it   passes 
through  along  with  the  metals  of  group  II.  (see  p.  274). 

IV.  To  the  fourth  portion  add  potassium  chromate.     Note 
the  yellow  precipitate  of  lead  chromate. 

Pb(NO3)2  +  K2Cr04  =  PbCrO4+  2KNO3. 

V.  To  the  fifth  portion  add  dilute  sulphuric  acid ;  a  white 
granular  precipitate  of  lead  sulphate  is  produced. 

Pb(NO3)2  +  H2SO4  =  PbS04  +  2HNO3. 

Dry  Reaction. — Powder  a  small  crystal  of  lead  nitrate, 
mix  about  twice  as  much  sodium  carbonate  with  it,  and  place 
the  mixture  in  a  shallow  cavity  scooped  out  on  a  piece  of 
charcoal.  Heat  the  mixture  by  means  of  a  blow-pipe  flame 
(see  Fig.  108,  p.  261),  holding  the  charcoal  in  such  a  position 
that  the  middle  or  inner  part  of  the  flame,  and  not  the  tip 
of  it,  plays  upon  the  mixture.  [The  outer  part  of  the  flame, 
where  the  oxygen  of  the  air  is  in  excess,  is  called  the  oxidizing 
flame ;  while  the  inner  portion,  where  heated  coal-gas  is  in 
excess,  is  known  as  the  reducing flame^\ 

The  mixture  quickly  melts,  and  the  lead  compound  is 
reduced  to  the  state  of  metallic  lead,  which  will  appear  in  the 
form  of  small  brilliant  globules  upon  the  charcoal.  At  the 
same  time,  some  of  the  lead  is  oxidized,  and  oxide  deposits 
round  the  cavity  as  a  yellow  incrustation.  Allow  the  mass  to 
cool ;  pick  out  one  of  the  globules  of  metal  with  a  penknife, 
and  show  that  it  is  soft  and  malleable  by  hammering  it ;  also 
that  if  rubbed  across  paper  it  leaves  a  black  mark. 
Copper,  Cu.  (Group  II.) 

Use  copper  sulphate,  CuSO4,5H2O.  Dissolve  a  little  of 
the  salt  in  water.  Note  that  it  is  readily  soluble. 

I.  Take  a  small  portion  of  the  solution  and  pass  sulphu- 
retted hydrogen  (the  group-reagent).  Notice  the  brownish- 
black  precipitate  of  copper  sulphide. 

CuSO4  -f  H2S  =  CuS  +  H2SO4. 

Filter  the  mixture,  and  show  that  the  precipitate  dissolves 
in  boiling  nitric  acid,  giving  a  bluish  solution.  (Compare  the 


Reactions  for  Iron.  277 

behaviour  of  lead  sulphide.)  Cautiously  add  ammonia  to 
this  solution,  and  note  the  deep  azure-blue  colour  of  the 
liquid. 

II.  To  a  second  portion  of  the  copper  sulphate  solution 
add  ammonia,  drop  by  drop.  Notice  the  pale  greenish-blue 
precipitate.  Add  more  ammonia,  and  the  precipitate  quickly 
dissolves,  forming  a  deep  blue  solution,  characteristic  of 
copper  compounds. 

Iron,  Fe.     (Group  1 110.) 

This  element  forms  two  classes  of  compounds,  which 
behave  quite  differently  towards  reagents.  One  of  these 
classes  is  derived  from  ferrous  oxide,  FeO,  and  the  other  from 
ferric  oxide,  Fe2O3,  in  which  the  iron  is  combined  with  a  larger 
proportion  of  oxygen,  or  is  in  a  higher  state  of  oxidation,  as  we 
say.  These  two  classes  of  iron  compounds  are  therefore 
distinguished  as  ferrous  and  ferric  salts.  The  former  are 
mostly  pale  green  (Exp.  51,  p.  46)  or  white,  while  the  ferric 
salts  are  generally  yellow. 

Ferrous  are  readily  converted  into  ferric  salts  by  the  action 
of  oxidizing  agents  ;  while  reducing  agents  change  ferric  back 
to  ferrous  compounds.  Thus,  if  sulphuretted  hydrogen  is  passed 
through  a  neutral  or  acid  solution  of  ferric  chloride,  ferrous 
chloride  is  produced  and  sulphur  is  precipitated. 

FeoCl6  +  H2S  =  2FeCl2  +  2HC1  +  S. 

Use  ferrous  sulphate,  FeSO4,7H2O,  and  ferric  chloride, 
Fe2Cl6. 

Make  a  solution  of  ferrous  sulphate  by  dissolving  two  or 
three  crystals  of  the  salt  in  cold  water. 

I.  Add  ammonia  (group-reagent),  and  obtain  a  dirty 
greenish  precipitate  of  ferrous  hydrate,  which  on  exposure  to 
the  air,  gradually  becomes  oxidized  into  ferric  hydrate  (brown). 

FeS04  +  2NH4HO  -  Fe(HO)2  +  (NH4)2SO4. 

Repeat  with  a  solution  of  ferric  chloride ;  a  reddish-brown 
precipitate  of  ferric  hydrate  is  formed. 

Fe2Cl6  -I-  6NH4HO  =  Fe2(HO)6  +  6NH4C1. 


278  Simple  Qualitative  Analysis. 

II.  Add  ammonium  sulphide  to   another  portion  of  the 
ferrous   sulphate   solution.      Black   ferrous  sulphide   is   pre- 
cipitated. 

FeSO4  +  (NH4)2S  =  FeS  +  (NH4)2SO4. 

Repeat  with  ferric  chloride.  The  same  precipitate  is 
obtained,  mixed  with  sulphur. 

Fe2Cl6  +  3(NH4)2S  =  2FeS  +  S  +  6NH4C1. 

III.  Add  a  few  drops  of  ammonium  thiocyanate  (frequently 
wrongly  called  ammonium  sulphocyanide)  to  the  ferrous  sul- 
phate.    No  change  takes  place. 

Repeat  with  ferric  chloride.  An  intense  blood-red  coloured 
solution  is  produced. 

IV.  Add  a   few  drops  of   potassium   ferri cyanide  to  the 
ferrous  sulphate.     A  dark  blue  precipitate  is  formed  (called 
Turnbull's  blue). 

Repeat  with  ferric  chloride.  No  blue  colour,  but  the 
mixture  becomes  brownish. 

V.  Add  potassium  ferrocyanide  to  ferrous  sulphate,  a  light 
blue  precipitate  is  formed,  which  on  exposure  to  air  becomes 
darker  blue. 

Repeat  with  ferric  chloride.  A  dark  blue  precipitate  re- 
sults (called  Prussian  blue). 

Zinc,  Zn.     (GROUP  III£.) 

Use  zinc  sulphate,  ZnSO4,7H2O.  Dissolve  a  few  crystals 
in  water.  Note  the  appearance  of  the  crystals  and  their  ready 
solubility. 

I.  Add  ammonium  sulphide  (group-reagent).     A  white 
precipitate  of  zinc  sulphide. 

ZnSO4  +  (NH4)2S  =  ZnS  +  (NH4)2SO4. 

II.  Add  ammonia  drop  by  drop  to  another  portion.     Note 
that  a  white  precipitate  is  at  first  produced ;  which,  however, 
readily  dissolves  as  more  ammonia  is  added. 

(This  reaction  enables  us  to  separate  zinc  from  iron.) 

III.  Dry  Reaction. — Heat  a  little  solid  zinc  sulphate 
with  sodium  carbonate  upon  charcoal  in  the  inner  blowpipe- 


Reactions  for  Calcium.  279 

flame.  No  metallic  beads  are  formed,  because  zinc  oxidizes 
too  easily  ;  but  an  incrustation  of  zinc  oxide  is  formed  on  the 
charcoal,  which  appears  canary  yellow  while  it  is  hot,  but 
turns  white  on  cooling.  Touch  the  incrustation  with  a  drop 
of  cobalt  nitrate  solution,  and  again  heat  it  in  the  extreme  tip 
of  the  flame.  T  he  mass  becomes  green. 

Calcium,  Ca.     (GROUP  IV.) 

Use  calcium  chloride,  CaCl2,6H2O.  Dissolve  some  of  the 
salt  in  water.  Note  how  quickly  and  easily  it  dissolves. 

I.  Add    ammonium   carbonate   (group-reagent)   to    a 
portion  of  the  solution.     A  white  precipitate  of  calcium  car- 
bonate is  obtained. 

CaCl2+  (NH4)2C03  -  CaC03  +  2NH4CL 

Filter  the  mixture,  and  pour  a  few  drops  of  acetic  acid  upon 
the  filter.  Note  effervescence  as  the  precipitate  dissolves. 

II.  Add  ammonium  oxalate  to  another  portion.     A  white 
precipitate  of  calcium  oxalate  is  obtained. 

CaCl2  +  (NH4)2C204  -  CaC204  +  2NH4C1. 

Pour  half  this  mixture  into  a  second  test-tube.  To  one 
portion  add  hydrochloric  acid.  Note  that  the  precipitate 
dissolves ;  to  the  other  add  acetic  acid.  The  precipitate  does 
not  dissolve. 

Dry  Reaction. — Dip  a  clean  platinum  wire  into  the 
calcium  chloride  solution,  and  bring  it  against  the  edge  of  a 
Bunsen  flame.  Note  the  reddish  colour  of  the  flame. 

Potassium,  K.     (GROUP  V.) 

Use  potassium  chloride,  KC1.  Dissolve  a  few  crystals  in 
a  small  quantity  of  water  so  as  to  obtain  a  strong  solution. 

I.  Add  platinum  chloride  (one  or  tw$  drops)  to  a  small 
quantity  of  this  solution.  A  yellow  crystalline  precipitate  is 
formed,  consisting  of  the  compound  K2PtCl6. 

Cautiously  add  water  drop  by  drop,  and  gently  warm  the 
mixture.  Notice  that  the  precipitate  dissolves.  Hence  this 
test  can  only  be  used  when  the  solution  is  strong.  Now  add 


280  Simple  Qualitative  Analysis. 

a  little  strong  hydrochloric  acid.  This  causes  the  reformation 
of  the  precipitate.  Therefore,  before  applying  this  test,  it  is 
best  to  add  a  drop  of  strong  hydrochloric  acid,  as  this  promotes 
the  formation  of  the  precipitate. 

II.  Add  to  a  second  portion  of  potassium  chloride  a  little 
strong  solution  of  hydrogen  sodium  tartrate,  a  white  precipitate 
of  hydrogen  potassium  tartrate  is  formed.  Add  a  little  water, 
and  note  that  the  precipitate  quickly  dissolves.  Therefore  this 
test  also  can  only  be  made  with  strong  solutions. 

Dry  Reactions. — Dip  a  clean  platinum  wire  into  the 
potassium  chloride  solution  and  bring  it  into  the  Bunsen  flame. 
A  lilac  colour  is  given  to  the  flame.  This  colour,  however, 
is  entirely  masked  by  a  yellow  colour  if  any  sodium  compounds 
are  present,  even  in  the  minutest  quantities.  If  the  flame  be 
looked  at  through  deep  blue  glass  (or  better  a  glass  prism 
filled  with  indigo)  the  colour  given  by  the  potassium  salt  will 
appear  crimson-red  in  spite  of  the  sodium  impurity. 

AMMONIUM  COMPOUNDS. 

These  all  evolve  ammonia,  when  heated  with  caustic  soda 
or  potash,  which  is  detected  by  its  smell,  and  by  its  action  on 
turmeric  or  litmus  paper.  See  p.  183. 

Reactions  for  Acids. — The  four  commonest  acids  are 
hydrochloric,  nitric,  carbonic,  and  sulphuric.  The  reactions 
by  which  these  are  distinguished  have  been  already  described. 
Chlorides,  p.  125  ;  Nitrates,  p.  195  ;  Carbonates,  p.  230,  and 
sulphates,  p.  261. 

Method  of  analysis  of  a  simple  mixture.1 

If  the  substance  given  is  solid,  its  general  appearance  (such 
as  its  colour,  whether  crystalline  or  not,  etc.)  should  be 
carefully  observed  aod  noted  down.  Then  proceed  to  make 
the  following  tests  directly  upon  the  solid. 

I.  Put  a  little  of  the  substance  into  a  short,  narrow  test- 
tube  and  heat  it.  Observe  closely  what  happens. 

1  Containing  only  chlorides,  nitrates,  carbonates,  or  sulphates  of 
ammonium,  potassium,  calcium,  zinc,  iron,  copper,  or  lead. 


Preliminary  Tests.  281 

(a)  If  a  white  sublimate  is  formed,  suspect  ammonium  salts. 
Confirm  this  by  heating  a  little  of  the  salt  with  caustic  soda. 

Ammonia  given  off,  detected  by  its  smell,  and  its  action  on 
turmeric  paper,  proves  the  presence  of  ammonium  salts. 

(b)  If  a  brown-coloured   vapour   is   given   off,  suspect   a 
nitrate,  probably  lead  nitrate. 

Confirm  a  nitrate  by  applying  the  test  for  nitric  acid,  given 
on  p.  195. 

II.  Heat  a  little  of  the  substance  (powdered)  on  charcoal 
in  the  blow- pipe  flame  (as  shown  on  p.  261). 

(a)  If  the  substance  deflagrates,  appearing  to  go  on  fire 
on  the  charcoal,  suspect  a  nitrate. 

(U)  If  a  bright  white  residue  is  left,  suspect  calcium  or  zinc 
compounds. 

Place  a  portion  of  the  residue  upon  turmeric  paper  and 
moisten  with  water.  If  a  brown  stain  is  produced,  suspect 
calcium. 

Moisten  a  part  of  the  residue  on  the  charcoal  with  cobalt 
nitrate,  and  again  heat.  If  a  green  residue  remains,  suspect 
zinc. 

(c)  If  metallic  beads  are  formed,  suspect  lead. 

Confirm  by  mixing  the  substance  with  sodium  carbonate, 
and  heating  on  charcoal  in  the  inner  blow-pipe  flame.  Mal- 
leable bead,  which  marks  paper,  confirms  lead. 

III.  Heat  a  little  of  the  substance  on  platinum  wire  in  a 
Bunsen  flame. 

(#).  If  reddish  flame,  suspect  calcium.  I  These  must  be  con- 
($)  If  lilac  flame,  suspect  potassium.  \        firmed  later. 

IV.  Add  a  little  dilute  hydrochloric  acid  to  some  of  the 
original  substance,  in  a  test-tube. 

(a)  If  effervescence  takes  place,  suspect  a  carbonate. 
Confirm  by  allowing  the  gas  to  enter  a  tube  containing 
lime  water  (as  shown  on  p.  227). 

V.  Gently  heat  a  little  of  the  substance  with  a  few  drops 
of  strong  sulphuric  acid. 

(a)  If  acid  fumes  are  evolved,  suspect  a  chloride  or  a  nitrate. 
Dip  a  glass  rod,  moistened  with  silver   nitrate,  into  the 
mouth  of  the  test-tube. 


282  Simple  Qualitative  Analysis. 

A  white  precipitate  on  the  rod  indicates  hydrochloric  acid. 

Drop  a  small  fragment  of  copper  into  the  tube.  Brown 
fumes  of  oxides  of  nitrogen  indicate  nitric  acid. 

(Hydrochloric  acid  may  also  be  confirmed  by  heating  the 
substance  with  manganese  dioxide  and  sulphuric  acid,  when 
chlorine  will  be  evolved,  which  is  recognized  by  its  colour 
and  its  bleaching  properties.) 

Prepare  a  solution  of  the  substance. — Place  a  little  of  the 
powder  in  a  test-tube,  add  water,  and  gently  warm  the  mixture. 
If  it  does  not  dissolve,  add  hydrochloric  acid,1  and  again  boil. 

When  the  solution  is  obtained,  proceed  according  to  the 
following  Table ;  except  that  if  hydrochloric  acid  has  been 
used  to  dissolve  the  substance  the  first  step  is  omitted. 

1  The  only  compounds  of  the  prescribed  metals  and  acids  (see  note, 
p.  280)  which  are  not  soluble  in  water  are  the  carbonates  of  lead,  copper, 
iron,  and  calcium  ;  and  the  sulphates  of  lead  and  calcium.  The  carbonates 
dissolve  readily  in  dilute  hydrochloric  acid,  with  evolution  of  carbon  dioxide : 
in  the  case  of  lead  carbonate,  lead  chloride  is  formed,  which  dissolves  on 
boiling,  but  crystallizes  out  again  on  cooling  (see  Reaction  I.,  p.  275). 
Lead  sulphate  and  calcium  sulphate  are  partially  dissolved  by  boiling 
water. 


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APPENDIX 

FOR  the  convenience  of  teachers  who  may  reside  in  the  country, 
and  who  have  to  order  chemicals  from  a  distance,  the  following 
list  has  been  drawn  up.  It  contains  the  substances  required  in  the 
carrying  out  of  the  experiments  described  in  this  book,  with  the 
quantities  which  it  would  be  suitable  to  obtain  for  a  small  class  of 
from  ten  to  twenty  students. 

The    approximate   cost,  including   the   bottles,  is  ,£3  Ss->  but 
exclusive  of  platinum  wire. 


i  Winchester  Ammonia  (sp.  g. 

0-880). 
.,  Hydrochloric 

acid  (pure). 
„  Nitric  acid  (1-42 

pure). 
„  Sulphuric       acid 

(pure  com.). 

i-lb.  bottle  Alcohol  (pure  abs.). 
„          Ammonium        sul- 
phide, 
lib.  Ammonium     chloride 

(pure). 
„     Lime. 

„     Manganese  dioxide. 
„     Marble. 

„     Potassium  chlorate. 
„  „         nitrate. 

„     Soda,  caustic,  sticks. 
„     Sodium  carbonate  (pure). 
„     Sulphur,  roll. 
„     Zinc,  granulated. 
I  Ib.  Alum. 

„     Ammonium  carbonate. 
„  „  nitrate. 

sulphate. 

„     Barium  chloride. 
„     Calcium  chloride  (cryst.). 
„    Copper  foil. 
„          „       gauze. 

„      sulphate. 
„     Ferrous  sulphate. 
„     Ferric  oxide. 
,,     Glycerin. 
„     Iron  pyrites. 
„        „    filings. 
„     Lead  acetate. 
„        „    dioxide. 
„         „     nitrate. 
„     Mercury. 
„     Mercuric  oxide  (red). 


lib. 


Oxalic  acid. 
Potash,  caustic. 

„     Potassium  chloride. 

„  „         dichromate. 

„     Red  lead. 

,,     Sodium  nitrate. 
„       nitrite. 

„     Sulphur,  flowers. 

„  Zinc  sulphate, 
i  oz.  Ammonium  thiocyanate. 

„     Aniline  blue  (or  magenta). 

„     Arsenious  oxide. 

„     Bromine. 

„     Calcium  carbide. 

„     Cobalt  chloride. 

„     Copper  chloride. 

„          „       nitrate. 

„          „       oxide. 

„     Ferric  chloride. 

„     Iodine. 

„     Litmus. 

„     Magnesium  carbonate. 

„     Mercuric  chloride. 

,,,     Phosphorus. 

„     Potassium  iodide. 

„  „         ferricyanide. 

„  „         ferrocyanide. 

„  „         permanganate. 

„     Sodium. 

„       oleate. 
„       peroxide. 

„  „       sulphite. 

„     Tartar  emetic. 

„  Tartaric  acid. 
£  oz.  Silver  nitrate. 
I  oz.  Magnesium. 

„     Potassium. 

i  drachm  Platinum  chloride, 
i  book  Gold  leaf. 
75  grains  Platinum  wire  (No.  29 
standard  wire  gauge). 


INDEX 


ABSOLUTE  zero,  102 

Atmosphere,  composition  of,  168 

Absorption  of  gases  by  charcoal,  215 

Atomic  symbols,  152 

Acetylene,  265 

weights,  148 

Acid,  acetic,  266 

Atoms,  149 

,  carbonic,  225 

Avogadro's  hypothesis,  162 

—  ,  formic,  265 

,  hydriodic,  134 

BALANCE,  92 

V»XT/1»-rtV\1-r\TV*l/^          T  1  A 

Barometer,  89 

"j  iiyuruuruiiiiCj  1,^4. 

,  hydrochloric,  120 

Bases,  69 

,  hydrofluoric,  134 

Basic  oxides,  63 

nifvir*       yQQ 

Bending  glass,  30 

,  mine,  ioo 

,  nitrous,  203 

Black-lead,  214 

,  oxalic,  233 

Bleaching  powder,  132 

,  stearic,  266 

Blue  vitriol,  82 

—  ,  sulphuric,  256 

Boiling  point  of  water,  80 

,  sulphurous,  253 

Bone  black,  217 

Acid-forming  oxides,  63 

Bordering  glass  tubes,  35 

Acids,  monobasic,  193 

Boring  corks,  33 

Affinity,  chemical,  II 

Boyle's  law,  98 

Alcohols,  266 

Erin's  process  (oxygen),  57 

Alkalies,  60,  66 

Bromine,  133 

Allotropy,  207 

Burette,  93 

Ammonia,  180 

Ammoniacal  liquor,  179 

CALORIE,  175 

Ammonium  chloride,  182 

Cane  sugar,  270 

nitrate,  182 

Carbohydrates,  270 

nitrite,  177 

Carbon,  210 

Analysis,  simple,  272 

A       f\£* 

Animal  charcoal,  217 

dioxide,  220 

Aqua  fortis,  192 

monoxide,  231 

Aqua  regia,  192 

Carbonates,  226 

Argon,  1  66 

Carbonic  acid,  22^ 

Atmosphere,  166 

Catalysis,  14 

286 


Index. 


Cavendish's  experiment,  72 
Chalk,  211 
Chamber  acid,  258 
Charcoal,  214 

,  absorption  of  gases  by,  215 

,  animal,  217 

Charles'  law,  97 
Chemical  action,  12 

,  modes  of,  155 

affinity,  1 1 

change,  6 

combination,  laws  of,  117 

equations,  154 

formulae,  153 

symbols,  152 

Chili  saltpetre,  180 
Chlorides,  tests  for,  126 
Chlorine,  128 
Coal,  217 
Coke,  216 
Collecting  gases,  24 
Combustion,  170 

,  gain  in  weight  by,  64 

,  heat  of,  174 

— ,  supporter  of,  52 

,  temperature  of,  172 

Compounds,  8 
Copper  nitrate,  16 

pyrites,  236 

,  tests  for,  276 

Crith,  101 

Critical  temperature,  103 
Crystallisation,  120 
Cubical  nitre,  194 

DECANTATION,  17 
Deliquescence,  83 
Densities  of  gases,  163 
Dephlogistigated  air,  55 
Destructive  distillation,  179 
Dextrin,  271 
Diamond,  213 
Diffusion  of  gases,  169 
Distillation,  23 
Dolomite,  211 
Double  decomposition,  156 


Downward  displacement,  27 
Dumas'  apparatus,  112 
Dutch  metal  in  chlorine,  131 

EFFLORESCENCE,  83 
Electrolysis  of  water,  76 
Elements  and  compounds,  8 
Equations,  chemical,  154 
Equivalents,  chemical,  139 
Ethylene,  212,  265 
Evaporation,  15 

FERMENTATION,  222 
Filtration,  17 
Fire  damp,  264 
Fitting  up  apparatus,  28 
Fluorine,  133 
Fluor-spar,  133 
Fusion,  21 

GALENA,  236 

Gases,  densities  of,  163 
— ,  diffusion  of,  169 
— ,  general  properties,  97 

,  relation  of  volume  to  heat,  97 

, •  pressure,  98 

Glass  blowing,  35 

Glass  tubes,  bending,  30 

Glucose,  270 

Glycerin,  268 

Graham's  law,  170 

Graphite,  214 

Green  vitriol,  46 

Gunpowder,  194 

Gypsum,  236 

HALOGENS,  133 
Hardness  of  water,  81 
Hydrocarbons,  264 
Hydrogen,  40 

chloride,  125 

. peroxide,  83 

Hydroxides,  63 

ICE,  79 

Ignition  point,  173 


Index. 


287 


Indestructibility  of  matter,  65 
Iodine,  134 
Iron,  tests  for,  277 
Iron  pyrites,  236 

LAMP-BLACK,  217 
Laughing  gas,  202 
Law  of  Boyle,  98 

Charles,  97 

constant  composition,  117 

Gay-Lussac,  162 

multiple  proportions,  118 

Lead,  tests  for,  275 
Lime-light,  173 
Limestone,  211 
Lime  water,  221 
Liquefaction  of  gases,  102 

MAGNESIUM  carbonate,  227 

combustion  in  air,  105 

oxide,  composition  of,  105 

Marble,  211,  228 
Marsh  gas,  210,  264 
Matter,  indestructibility  of,  65 
Mechanical  mixtures,  9 
Mercuric  oxide,  54,  115 
Metalloids,  9,  178 
Metals  and  non-metals,  9 
Metric  system,  85 
Molecular  formulae,  153 

weights  of  gases,  163 

Molecules,  149 
Monobasic  acids,  193 
Mortar,  228 

NASCENT  state,  130 
Natural  waters,  81 
Neutralization,  68,  143 
Nitrates,  detection  of,  195 
Nitre,  10,  180 
Nitric  oxide,  197 
Nitrites,  203 
Nitrogen,  176 

— ,  oxides  of,  197 
Nitrous  oxide,  201 


OLEFIANT  gas,  265 

Oxidation,  63 

Oxides,  62 

Oxygen,  54 

Oxyhydrogen  flame,  173 

Ozone,  206 

,  tests  for,  208 

PERCENTAGE  composition,  114 
Permanent  hardness  of  water,  82 
Phosphorus,  combustion  in  oxygen, 

59 

Pipette,  93 
Plaster  of  Paris,  224 
Plumbago,  214 
Potassium,  tests  for,  279 

permanganate,  203 

Precipitation,  19 
Prussian  blue,  193 
Pyrites  burners,  257 

QUANTITATIVE  chemical  equations, 

158 

Quartz,  214 
Quicksilver,  2 

RAIN  water,  81 
Reagents,  273 
Red  lead,  218 
Respiration,  65 
Rock  crystal,  214,  219 
Rust,  64 

SAL  ammoniac,  4,  182 

Saltpetre,  180 

Salts,  69 

Saturated  solutions,  20 

Silver  displaced  by  zinc,  140 

Simple  manipulations,  15 

Soap,  268 

Soap-bubbles  with  hydrogen,  50 

Sodium,  action  on  water,  42 

carbonate,  227 

peroxide,  56 

Solubility  of  gases  in  water,  81 
of  salts  in  water,  2 1 


288 


Index. 


Solution,  15 

Solvent  power  of  water,  So 
Spirits  of  hartshorn,  180 
Starch,  211,  270 

compound  with  iodine,  134 

States  of  matter,  i 
Steam,  4 

Sugar  of  lead,  216 
Sublimation,  5 
Sulphates,  259 
Sulphides,  241,  247 
Sulphites,  253 
Sulphur,  236 

,  allotropic  forms  of,  239 

dioxide,  250 

Sulphuretted  hydrogen,  243 

TEMPORARY  hardness,  water,  82 
Thermal  unit,  175 
Thermometers,  87 
Torricellian  vacuum,  90 

UNIT  volume,  164 
Upward  displacement,  27 


VITRIOL  chambers,  257 
Vitriols,  259 

WASHING  soda,  224 
Water,  72 

,  colour  of,  78 

,  electrolysis  of,  76 

,  freezing  of,  80 

,  gaseous,  4 

,  gravimetric    composition  of, 

112 

,  hardness  of,  81 

,  maximum  density  of,  79 

,  of  crystallization,  82 

,  sea,  8 1 

,  volume  composition  of,  77 

Water-gas,  232 
Waters,  natural,  81 
Weighing  and  measuring,  85 

ZINC  blende,  236 

,  granulated,  45 

,  tests  for,  278 


PRINTED   BY  WILLIAM   CLOWES   AND  SONS,   LIMITED, 
LONDON   AND   BECCLES. 


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